317 terms

Chem Ch 2

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Properties of Waves
To understand electromagnetic radiation and be able to perform calculations involving wavelength, frequency, and energy.
Several properties are used to define waves. Every wave has a wavelength, which is the distance from peak to peak or trough to trough. Wavelength, typically given the symbol λ (lowercase Greek "lambda"), is usually measured in meters. Every wave also has a frequency, which is the number of wavelengths that pass a certain point during a given period of time. Frequency, given the symbol ν (lowercase Greek "nu"), is usually measured in inverse seconds (s−1). Hertz (Hz), another unit of frequency, is equivalent to inverse seconds.
The product of wavelength and frequency is the speed in meters per second (m/s). For light waves, the speed is constant. The speed of light is symbolized by the letter c and is always equal to 2.998×108 m/s in a vacuum; that is,
c=λν=2.998×108m/s
Another term for "light" is electromagnetic radiation, which encompasses not only visible light but also gamma rays, X-rays, UV rays, infrared rays, microwaves, and radio waves. As you could probably guess, these different kinds of radiation are associated with different energy regimes. Gamma rays have the greatest energy, whereas radio waves have the least energy. The energy (measured in joules) of a photon for a particular kind of light wave is equal to its frequency times a constant called Planck's constant, symbolized h:
Ephoton=hν
where
h=6.626×10−34J⋅s
These two equations can be combined to give an equation that relates energy to wavelength:
E=hcλ
A radio station's channel, such as 100.7 FM or 92.3 FM, is actually its frequency in megahertz (MHz), where 1MHz=106Hz and 1Hz=1s−1.
Calculate the broadcast wavelength of the radio station 93.10 FM.
λ =
3.220 m
Green light has a frequency of about 6.00×1014s−1. What is the energy of a photon of green light?
Ephoton =
3.98×10−19 J
Hospital X-ray generators emit X-rays with wavelength of about 15.0 nanometers (nm), where 1nm=10−9m. What is the energy of a photon of the X-rays?
Ephoton =
1.33×10−17 J
Which type of electromagnetic radiation has the greatest energy?
Gamma rays
What is the wavelength in meters of an FM radio wave with frequency ν = 149.0 MHz ?
λFM =
2.01
m
What is the wavelength of a medical X ray with ν = 5.55×1017 Hz ?
λXray =
5.41×10−10
m
What is the frequency of a gamma ray with 5.17×10−11 m ?
νgamma =
5.80×1018
Hz
What is the frequency of a radar wave with 17.3 cm ?
1.73×109
Hz
The Rydberg Equation
An astrophysicist working at an observatory is interested in finding clouds of hydrogen in the galaxy. Usually hydrogen is detected by looking for the Balmer series of spectral lines in the visible spectrum. Unfortunately, the instrument that detects hydrogen emission spectra at this particular observatory is not working very well and only detects spectra in the infrared region of electromagnetic radiation. Therefore the astrophysicist decides to check for hydrogen by looking at the Paschen series, which produces spectral lines in the infrared part of the spectrum. The Paschen series describes the wavelengths of light emitted by the decay of electrons from higher orbits to the n=3 level.
What wavelength λ should the astrophysicist look for to detect a transition of an electron from the n=5 to the n=3 level?
λ =
1.28×10−6 m
The Photoelectric Effect
Electrons are emitted from the surface of a metal when it's exposed to light. This is called the photoelectric effect. Each metal has a certain threshold frequency of light, below which nothing happens. Right at this threshold frequency, an electron is emitted. Above this frequency, the electron is emitted and the extra energy is transferred to the electron.
The equation for this phenomenon is
KE=hν−hν0
where KE is the kinetic energy of the emitted electron, h=6.63×10−34J⋅s is Planck's constant, ν is the frequency of the light, and ν0 is the threshold frequency of the metal.
Also, since E=hν, the equation can also be written as
KE=E−E0
where E is the energy of the light and E0 is the threshold energy of the metal.
Here are some data collected on a sample of cesium exposed to various energies of light.
What is the threshold frequency ν0 of cesium?
Note that 1 eV (electron volt)=1.60×10−19 J.
ν0 =
9.39×1014
Hz
A ray of red light has a wavelength of about 7.0×10−7 m. Will exposure to red light cause electrons to be emitted from cesium?
no
What is the kinetic energy of the emitted electrons when cesium is exposed to UV rays of frequency 1.9×1015Hz?
KE =
6.37×10−19
J
...
Electromagnetic radiation behaves both as particles (called photons) and as waves. Wavelength (λ) and frequency (ν) are related according to the equation
c=λ×ν
where c is the speed of light (3.00×10^8 m/s). The energy (E in joules) contained in one quantum of electromagnetic radiation is described by the equation
E=h×ν
where h is Planck's constant (6.626×10^−34 J⋅s). Note that frequency has units of inverse seconds (s^−1), which are more commonly expressed as hertz (Hz).
A microwave oven operates at 2.30 GHz . What is the wavelength of the radiation produced by this appliance?
λ =
1.30×108
nm

Some people lose their wireless Internet connection at home while their microwave oven is turned on because both happen to operate near 2.40 GHz.
Two of the types of ultraviolet light, UVA and UVB, are both components of sunlight. Their wavelengths range from 320 to 400 nm for UVA and from 290 to 320 nm for UVB. Compare the energy of microwaves, UVA, and UVB.
UVB radiation causes sunburn whereas UVA radiation does not. However, UVA, which causes tanning, is thought to be even more dangerous. The precise wavelengths of ultraviolet light that contribute to the formation of skin cancers still need to be determined by scientists.
Emission Line Energy
The Rydberg equation expresses the wavelength, λ, of emitted light based on the initial and final energy states (ni and nf) of an electron in a hydrogen atom:
With some manipulation, the Rydberg equation can be rewritten in the form
E=constant×((1/nf²)−(1/ni²))
which allows you to calculate the energy of the emitted light. What is the value of the constant needed to complete this equation?
2.18×10−18
J
What is the change in energy, ΔE, in kilojoules per mole of hydrogen atoms for an electron transition from n=9 to n=2?
ΔE =
-312
kJ/mol

Each hydrogen atom emits energy in the form of a photon of light and hence the energy carries a negative sign. The emission spectra of various elements are used for lighting purposes as in sodium and mercury vapor lights as well as neon signs.
The Bohr Equation
The electron from a hydrogen atom drops from an excited state into the ground state. When an electron drops into a lower-energy orbital, energy is released in the form of electromagnetic radiation.
How much energy does the electron have initially in the n=4 excited state?
En =
−1.37×10−19 J
What is the change in energy if the electron from Part A now drops to the ground state?
ΔE =
−2.05×10−18 J

Energy was released in this transition, so we express ΔE as a negative number (it is a net loss of energy from the point of view of the system). However, you should use the absolute value of ΔE for the remaining calculations.
What is the wavelength λ of the photon that has been released in Part B?
λ =
9.70×10−8 m
What might the photon from Part C (slide above) be useful for?
Getting a suntan
For the electronic transition from n = 3 to n = 8 in the hydrogen atom, calculate the energy.
2.08 × 10−19 J
What is the longest-wavelength line in nanometers in the infrared series for hydrogen where m = 3?
λ =
1875
nm
What is the shortest-wavelength line in nanometers in the infrared series for hydrogen where m = 3?
λ =
820.4
nm
The biological effects of a given dose of electromagnetic energy generally become more serious as the energy of the radiation increases: Infrared radiation has a pleasant warming effect; ultraviolet radiation causes tanning and burning; and X rays can cause considerable tissue damage. What energies in kilojoules per mole are associated with the following wavelengths:

infrared radiation with λ = 1.62×10−6 m ?
73.9
kJ/mol
The biological effects of a given dose of electromagnetic energy generally become more serious as the energy of the radiation increases: Infrared radiation has a pleasant warming effect; ultraviolet radiation causes tanning and burning; and X rays can cause considerable tissue damage. What energies in kilojoules per mole are associated with the following wavelengths:

ultraviolet light with λ = 231 nm ?
518
kJ/mol
The biological effects of a given dose of electromagnetic energy generally become more serious as the energy of the radiation increases: Infrared radiation has a pleasant warming effect; ultraviolet radiation causes tanning and burning; and X rays can cause considerable tissue damage. What energies in kilojoules per mole are associated with the following wavelengths:

X rays with λ = 4.73 nm ?
2.53×104
kJ/mol
What is the energy in kilojoules per mole of photons corresponding to the shortest-wavelength line in the series of the hydrogen spectrum when m=1 and n>1 ?
E =
1310
kJ/mol
How many photons of frequency 1.50 × 1014 s−1 are needed to give 30.1 J of energy?
3.03 × 10^20 photons
The work function of iron metal is 451 kJ/mol. Will photons of violet light with = 390 nm cause electrons to be ejected from a sample of iron?
Photons with λ = 390 nm won't eject electrons from a sample.
As an electron drops from the n=5 level to the n=2 level, ____.
light of one color is emitted
Calculate in kilojoules per mole the energy necessary to completely remove an electron from the first shell of a hydrogen atom (R∞ = 1.097×10^−2nm^−1).
E =
1310
kJ/mol
What is the de Broglie wavelength of a 1.22 × 106 g truck that is moving at 105 km/hour ?
1.86 × 10−38m
What is the de Broglie wavelength in meters of a small car with a mass of 1150 kg traveling at a velocity of 55.0 mi/h (24.6 m/s)?
λ =
2.34×10^−38
m
Is this wavelength longer or shorter than the diameter of an atom (approximately 200 pm)?
shorter
The Heisenberg Uncertainty Principle
A student is examining a bacterium under the microscope. The E. coli bacterial cell has a mass of m = 1.60 fg (where a femtogram, fg, is 10−15g) and is swimming at a velocity of v = 8.00 μm/s , with an uncertainty in the velocity of 9.00 % . E. coli bacterial cells are around 1 μm ( 10−6 m) in length. The student is supposed to observe the bacterium and make a drawing. However, the student, having just learned about the Heisenberg uncertainty principle in physics class, complains that she cannot make the drawing. She claims that the uncertainty of the bacterium's position is greater than the microscope's viewing field, and the bacterium is thus impossible to locate.
What is the uncertainty of the position of the bacterium?
Δx =
4.58×10^−11 m
By looking at the uncertainty of the bacterium's position, did the student have a valid point?
The student is wrong. The uncertainty of the bacterium's position is tiny compared to the size of the bacterium itself.
Quantum Number Rules
Quantum numbers can be thought of as labels for an electron. Every electron in an atom has a unique set of four quantum numbers.
The principal quantum number n corresponds to the shell in which the electron is located. Thus n can therefore be any integer. For example, an electron in the 2p subshell has a principal quantum number of n=2 because 2p is in the second shell.
The azimuthal or angular momentum quantum number ℓ corresponds to the subshell in which the electron is located. s subshells are coded as 0, p subshells as 1, d as 2, and f as 3. For example, an electron in the 2p subshell has ℓ=1. As a rule, ℓ can have integer values ranging from 0 to n−1.
The magnetic quantum number mℓ corresponds to the orbital in which the electron is located. Instead of 2px, 2py, and 2pz, the three 2p orbitals can be labeled −1, 0, and 1, but not necessarily respectively. As a rule, mℓ can have integer values ranging from −ℓ to +ℓ.
The spin quantum number ms corresponds to the spin of the electron in the orbital. A value of 1/2 means an "up" spin, whereas −1/2 means a "down" spin.
What is the only possible value of mℓ for an electron in an s orbital?
0

Since the allowed values for mℓ range from −ℓ to +ℓ, once you know the value for ℓ you know the values for mℓ.
What are the possible values of mℓ for an electron in a d orbital?
-2,-1,0,1,2

Since the allowed values for mℓ range from −ℓ to +ℓ, once you know the value for ℓ you know the values for mℓ.
Which of the following set of quantum numbers (ordered n, ℓ, mℓ, ms) are possible for an electron in an atom?
4, 2, -1, -1/2
5, 2, 1, -1/2
4, 2, 1, -1/2

The allowed values for mℓ range from −ℓ to +ℓ and the allowed values for ℓ are integers between zero and n−1. Once you know the value for n, you can determine the acceptable ℓ values, and from there the acceptable mℓ values. The ms values are fixed at either 1/2 or −1/2.
Quantum Numbers
Every electron in an atom is described by a unique set of four quantum numbers: n, ℓ, mℓ, and ms. The principal quantum number, n, identifies the shell in which the electron is found. The angular momentum quantum number, ℓ, indicates the kind of subshell. The magnetic quantum number, mℓ, distinguishes the orbitals within a subshell. The spin quantum number, ms, specifies the electron spin.
Identify which sets of quantum numbers are valid for an electron. Each set is ordered (n,ℓ,mℓ,ms).
1,0,0,1/2
4,3,2,1/2
2,1,1,1/2
3,2,2,1/2
2,1,-1,1/2
Identify the sets of quantum numbers that describe all the electrons in the ground state of a neutral beryllium atom, Be. Each set is ordered (n,ℓ,mℓ,ms).
How many electrons can an n = 6 shell theoretically hold?
72 electrons
Give the possible combinations of quantum numbers for the following orbitals.

A 3s orbital
n = 3, l = 0, ml= 0
Give the possible combinations of quantum numbers for the following orbitals.

A 2p orbital
n = 2, l = 1, ml= -1,0,1
Give the possible combinations of quantum numbers for the following orbitals.

A 4d orbital
n = 4, l = 2, ml= -2,-1,0,1,2
Characteristics of an Atomic Orbital
Wave functions provide information about an electron's probable location in space. This can be represented by an electron-density distribution diagram, called an orbital. An orbital is characterized by a size, shape, and orientation in space.
What is the azimuthal quantum number (also called the angular-momentum quantum number), ℓ, for the orbital shown here?
ℓ= 2

For the known elements, only s, p, d, and f orbitals are used. However, quantum theory predicts the existence of orbitals with values higher than ℓ=3. For example, an orbital with ℓ=4 would be given the letter designation of g.
What is the label for the orbital shown here that indicates the type of orbital and its orientation in space?
dxz
Compare the orbital shown in Parts A and B (the two previous slides) to the orbital shown here in size, shape, and orientation.

Which quantum number(s) would be different for these two orbitals?
mℓ only

The label for this orbital would be dx2−y2.
The actual value of mℓ assigned to a given orientation is not arbitrary. It is determined based on how the hydrogen atom behaves in a magnetic field. This also accounts for the name given to this quantum number.
How would the dx2−y2 orbital in the n=5 shell compare to the dx2−y2 orbital in the n=3 subshell?
The contour of the orbital would extend further out along the x and y axes.
The value of ℓ would increase by 2.
The radial probability function would include two more nodes.
The orientation of the orbital would be rotated 45∘ along the xy plane.
The mℓ value would be the same.
True: A, C, E

False: B, D

The following representation of this orbital, 5dx2−y2, depicted when it is bisected by the xy plane, shows the effect of the radial nodes on the orbital contours.
What is the maximum number of electrons that a d subshell can hold?
10 electrons
Give a possible combination of n and l quantum numbers for the following fourth-shell orbital:
n=4, l=2
In the animation, you can see that the electrons occupy different orbitals according to the energy level of each orbital. A single box represents an orbital. The unpaired electron is represented assingle harpoon upwhereas the paired electrons in the same orbital are represented by two arrows pointing in opposite directions:single harpoon up and single harpoon down. Watch the animation and identify which of the following statements are correct.
The C atom has two unpaired electrons.

In the Li atom, the 3s, 3p, and 3d orbitals have different energies.

Electrons generally occupy the lowest energy orbital first.

The arrangement of the orbitals in a multielectron atom is different from the arrangement in a single-electron atom owing to the electron-electron repulsions in a multielectron atom. In the case of a single-electron atom, the orbitals in a given principal shell have the same energy. However, in the case of a multielectron atom, the orbitals in a given principal shell have different energies. Electrons occupy the lowest energy orbital first. Each orbital can hold a maximum of two electrons of opposite spins. When more than one orbital of equal energy is available, electrons will first occupy these orbitals singly with parallel spins.
Thus, the C atom has two unpaired electrons in its 2p subshell. The element that follows C is N. It has three unpaired electrons in the 2p subshell. In the N atom, all the three degenerate 2p orbitals are filled with single-electrons each. Thus, it has attained half-filled orbitals. For the next atom, oxygen, the pairing of electrons will occur.

The filling of the electrons in the different orbitals of an atom determines the electron configuration of the atom and indicates the presence or absence of unpaired electrons in the atom.
Rules for writing electron configuration
The electron configuration of an atom describes how the electrons fill the orbitals within an atom. Two of the rules that explain how electrons fill orbitals are as follows:
Hund's rule of maximum multiplicity states that when more than one orbital of equal energy is available, electrons will first occupy these orbitals singly with parallel spins. The pairing of electrons will start only after all the degenerate orbitals are singly occupied or are half-filled.
Pauli's exclusion principle states that each orbital can hold a maximum of two electrons of opposite spins.
For example, the electron configuration of a C atom with the atomic number 6 is (see picture)
Here, a single box represents an orbital, and an electron is represented as a half arrow. Orbitals of equal energy are grouped together.
According to Pauli's exclusion principle, each orbital can hold a maximum of two electrons of opposite spins. If you observe the electron configuration of the carbon atom, 1s and 2s orbitals hold two electrons of opposite spins. The fifth and sixth electrons enter the 2p orbital.
Because the 2p subshell has three orbitals of equal energy, according to Hund's rule, the fifth and sixth electrons occupy 2p orbitals singly with parallel spins instead of pairing.
Consider that a single box represents an orbital, and an electron is represented as a half arrow. Orbitals of equal energy are grouped together. The electron configuration will be valid if it follows both Hund's rule and Pauli's principle. Sort the various electron configurations based on whether they are valid, whether they break Hund's rule, or whether they break Pauli's exclusion principle.
The electron configuration can also be represented by writing the symbol for the occupied subshell and adding a superscript to indicate the number of electrons in that subshell. For example, consider a carbon atom having an atomic number of 6. The total number of electrons in a neutral carbon atom is 6. The electron configuration of the carbon atom represented by the orbital diagram is (see picture)
This electron configuration can be written as
1s22s22p2
where 1s, 2s, and 2p are the occupied subshells, and the superscript "2" is the number of electrons in each of these subshells.
Use the rules for determining electron configurations to write the electron configuration for Si.

1s^22s^22p^63s^23p²
The electron configuration for Si is 1s22s22p63s23p2.
From this method of writing the electron configuration, you cannot predict the number of unpaired electrons. You can write the electron configuration as an orbital diagram as (see picture)...
...where each orbital is denoted by a box and each electron is denoted by a half arrow.
Anomalous electron configurations
Some atoms, such as some transition metals and some elements in the lanthanide and actinide series, do not adhere strictly to Hund's rule and Pauli's principle. The reason the anomalies are observed is the unusual stability of both half-filled and completely filled subshells.
This behavior can be explained with an example of the chromium atom. Using Hund's rule and Pauli's principle, you can write the expected electron configuration of the Cr atom that strictly follows these rules as 1s22s22p63s23p64s23d4 . However, by moving an electron from the 4s orbital to the 3d orbital you obtain a half-filled 3d orbital. This half-filled orbital is more stable than the combination of the filled 4s orbital and the partially filled 3d orbital. Thus, the observed electron configuration of the Cr atom is 1s22s22p63s23p64s13d5.
Cu has an anomalous electron configuration. Write the observed electron configuration of Cu.
1s^22s^22p^63s^23p^64s^13d^10
Electron Configurations of Atoms
An atom consists of a small, positively charged nucleus, surrounded by negatively charged electrons. We organize the electrons in a logical manner. As the atomic number increases, electrons are added to the subshells according to their energy. Lower energy subshells fill before higher energy subshells.
The order of filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
The periodic table can be used to help you remember the order.
Give the complete ground-state electron configuration for silicon (Si).
1s^22s^22p^63s^23p^2
Give the ground-state electron configuration for silicon (Si) using noble-gas shorthand.
[Ne]3s^23p^2
Give the actual ground-state electron configuration for copper (Cu) using the complete form.
1s^22s^22p^63s^23p^64s^13d^10

The expected ground-state electron configuration of copper is 1s22s22p63s23p64s23d9; however, the actual configuration is 1s22s22p63s23p64s13d10 because a full d subshell is particularly stable. There are 18 other anomalous elements for which the actual electron configuration is not what would be expected.
Give the ground-state electron configuration for copper (Cu) using noble-gas shorthand.
[Ar]4s^13d^10
Electron Configurations: Rules and Principles
When writing the ground-state electron configuration of a many-electron atom, three main rules must be followed:
The aufbau principle: Electrons are added to the lowest energy orbitals available.
The Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers (n, ℓ, mℓ, and ms).
Hund's rule: For degenerate orbitals, the lowest energy state is attained when the number of electrons with the same spin is maximized. So for a degenerate set of orbitals, one electron goes into each orbital until all the orbitals of the subshell are half-filled. Once all the orbitals of the subshell are half-filled the pairing of electrons can take place.
Note that aufbau is the German word for "building up."
Classify each orbital diagram for ground-state electron configurations by the rule or principle it violates.
The following sets of quantum numbers, listed in the order n, ℓ, mℓ, and ms, were written for the last electrons added to an atom. Identify which sets are valid and classify the others by the rule or principle that is violated.
Orbital-Filling Diagrams
An orbital-filling diagram shows the number of electrons in each orbital, which are shown in order of energy. The placement of electrons in orbitals follows a certain set of rules.
Lower energy subshells fill before higher energy subshells. The order of filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. The periodic table can be used to help you remember this order.
An orbital can hold up to two electrons, which must have opposite spins.
Hund's rule states that if two or more orbitals with the same energy are available, one electron goes in each until all are half full. The electrons in the half-filled orbitals all have the same value of their spin quantum number.
How many orbitals are there in the third shell (n=3)?
9

Nine orbitals (one s, three p, and five d) can hold a maximum of 18 electrons.
Show the orbital-filling diagram for N (nitrogen). Stack the subshells in order of energy, with the lowest-energy subshell at the bottom and the highest-energy subshell at the top.
Show the orbital-filling diagram for S (sulfur). Stack the subshells in order of energy, with the lowest-energy subshell at the bottom and the highest-energy subshell at the top.
Show the orbital-filling diagram for Br (bromine). Stack the subshells in order of energy, with the lowest-energy subshell at the bottom and the highest-energy subshell at the top.
Identify the specific element that corresponds to the following electron configuration:
[Kr]5s 24d 105p 4.
Te
Relating Quantum Numbers and Electron Configurations to the Periodic Table
The periodic table lists all known elements arranged by atomic number. Atomic number is the nuclear charge, the number of protons in the nucleus of an an atom of a particular element. For a neutral atom, the number of protons is equal to the number of electrons.
Each column of the table, called a group, contains elements with the same number of valence electrons that are in different quantum levels. Each row of the table, called a period, contains elements with differing numbers of valence electrons that are in the same principal quantum level.
The four main blocks of the table (s, p, d, and f) contain elements whose highest energy electrons have the same azimuthal quantum number (ℓ).
For each set of elements represented in this periodic table outline, identify the principal quantum number, n, and the azimuthal quantum number, ℓ, for the highest energy electrons in an atom of one of those elements.
For an electron in an atom, specifying that n=2, ℓ=1 is equivalent to saying that it is located in the 2p subshell. Similarly, n=3, ℓ=0 specifies the 3s subshell, and so forth. Thus, one can construct the following table:
Outer electron configurations
The outer electron configuration of an element includes everything except the noble-gas core. For example, the element C has an electron configuration of [He]2s22p2 and an outer electron configuration of 2s22p2. Similarly, the element Pb has an electron configuration of [Xe]6s24f145d106p2 and an outer electron configuration of 6s24f145d106p2.
The black line between elements 56 and 71 in the periodic table shown indicates that in the Lanthanide series elements 57 through 70 are listed below the main table, while in the Actinide series elements 89-102 are listed below the main table. Elements 71 and 103 are listed in main table.
Identify the outer electron configuration of each element shown in this periodic table outline.
Identify the general outer electron configuration for each group of elements shown in this periodic table outline.
In what period and group on the periodic table, respectively, is the element with the following electron configuration located and what is the element's identity?
[Kr]5s 24d 105p 4
Period 5, group 6A, tellurium (Te)
Identify the atom with the following ground-state electron configuration:
Fe
Give expected ground-state electron configurations for the following atoms.
Ti (Z = 22)
1s^22s^22p^63s^23p^64s^23d^2
Give expected ground-state electron configurations for the following atoms.
Zn (Z = 30)
1s^22s^22p^63s^23p^64s^23d^10
Give expected ground-state electron configurations for the following atoms.
Sn (Z = 50)
1s^22s^22p^63s^23p^64s^23d^104p^65s^24d^105p^2
Give expected ground-state electron configurations for the following atoms.
Pb (Z = 82)
1s^22s^22p^63s^23p^64s^23d^104p^65s^24d^105p^66s^24f^145d^106p^2
Take a guess. What do you think is a likely ground-state electron configuration for the sodium ion, Na+, formed by loss of an electron from a neutral sodium atom?
What is a likely ground-state electron configuration for the chloride ion, Cl−, formed by adding an electron to a neutral chlorine atom?
Electron configurations which of the following elements are anomalous?
Th
Mo
Cu
Ag
U
Pt
Each element in the periodic table has a distinctive atomic radius.
Place the following elements in order of decreasing atomic size: arsenic, sulfur, neon, cesium, calcium, and phosphorus.
Atomic Radii and Effective Nuclear Charge
The atomic radius of an element can be predicted based on its periodic properties. Atomic radii increase going down a group in the periodic table, because successively larger valence-shell orbitals are occupied by electrons. Atomic radii generally decrease moving from left to right across a period because the effective nuclear charge increases.
Rank the following elements in order of decreasing atomic radius.
Atomic radii increase going down a group, because successively larger valence-shell orbitals are occupied by electrons. For example, rubidium has electrons in the fifth shell, which contains much larger orbitals than the fourth, third, second, or first shells.
Rank the following elements in order of decreasing atomic radius.
The shielding of electrons gives rise to an effective nuclear charge, Zeff, which explains why boron is larger than oxygen. Estimate the approximate Zeff felt by a valence electron of boron and oxygen, respectively?
+3 and +6

The valence electrons in an oxygen atom are attracted to the nucleus by a positive charge nearly double that of boron. Therefore, the electrons in oxygen are held closer to the nucleus, giving it a smaller radius.
When you compare the atomic radius of silicon (Si) to that of phosphorus (P), ____.
Silicon is a larger atom than phosphorus because phosphorus has one more proton than silicon, thereby increasing the attraction for all electrons and decreasing the atomic radius.
Which atom in the following pair would you expect to be larger? Ba or Mg. Explain.
Ba, Atoms get larger as you go down a group.
Which atom in the following pair would you expect to be larger?Rh or Nb. Explain.
Nb. Atoms get smaller as you go across a period.
Which atom in the following pair would you expect to be larger? S or Te. Explain.
Te. Atoms get larger as you go down a group.
Which atom in the following pair would you expect to be larger? Os or Lu. Explain.
Lu. Atoms get smaller as you go across a period.
Why do atomic radii increase going down a group of the periodic table?
Atomic radii increase going down a group because the electron shells are farther away from the nucleus.
Why do atomic radii decrease from left to right across a period of the periodic table?
Atomic radii decrease from left to right across a period because the effective nuclear charge increases.
Order the following atoms according to increasing atomic radius: S, F, O.
Which atom in each of the following pairs has a larger radius?
Na or K
K
Which atom in each of the following pairs has a larger radius?
V or Ta
Ta
Which atom in each of the following pairs has a larger radius?
V or Zn
V
Which atom in each of the following pairs has a larger radius?
Li or Ba
Ba
Where on the blank outline of the periodic table do elements that meet the following descriptions appear?

a. Elements with the valence-shell ground-state electron configuration ns2np5

b. An element whose fourth shell contains two p electrons

c. An element with the ground-state electron configuration [Ar]4s2 3d10 4p5
a. 7a group

b. Ge

c. Br
What atom has the following orbital-filling diagram (Figure 1) ?
Ga
Which of the following three spheres (Figure 1) represents a Ca atom, which an Sr atom, and which a Br atom?
Sr(215pm)>Ca(197pm)>Br(114pm)
Which has the higher frequency, red light or violet light?
violet light
Which has the longer wavelength, red light or violet light?
red light
Which has the greater energy, red light or violet light?
violet light
The Hubble Space Telescope detects electromagnetic energy in the wavelength range 1.15×10^−7m to 2.0×10^−6m.

a. What region of the electromagnetic spectrum is found completely within this range?

b. What regions fall partially in this range?
a. visible

b. ultraviolet and infrared
What is the wavelength (in meters) of ultraviolet light with ν = 5.4×10^15 s−1 ?
λ =
5.6×10^−8
m
A certain cellular telephone transmits at a frequency of 910 MHz and receives at a frequency of 955 MHz.

a. What is the wavelength of the transmitted signal in cm?

b. What is the wavelength of the received signal in cm?
a. 33.0 cm

b. 31.4 cm
The data encoded on CDs, DVDs, and Blu-ray discs is read by lasers.

a. What is the wavelength in nanometers of the CD laser (ν = 3.85×10^14 s−1)?

b. What is the energy in joules of the CD laser (ν = 3.85×10^14 s−1)?

c. What is the wavelength in nanometers of the DVD laser (ν= 4.62×10^14 s−1)?

d. What is the energy in joules of the DVD laser (ν= 4.62×10^14 s−1)?

e. What is the wavelength in nanometers of the Blu-ray laser ( ν= 7.41×10^14 s−1)?

f. What is the energy in joules of the Blu-ray laser ( ν= 7.41×10^14 s−1)?
a. 779 nm

b. 2.55×10^−19 J

c. 649 nm

d. 3.06×10^−19 J

e. 405 nm

f. 4.91×10^−19 J
According to the equation for the Balmer line spectrum of hydrogen, a value of n = 3 gives a red spectral line at 656.3 nm, a value of n = 4 gives a green line at 486.1 nm, and a value of n = 5 gives a blue line at 434.0 nm.

Calculate the energy in kilojoules per mole of the radiation corresponding to each of these spectral lines.
E3, E4, E5 =
182.3,246.1,275.6
kJ/mol
The work function of silver metal is 436 kJ/mol. What frequency of light is needed to eject electrons from a sample of silver?
1.09×10^15 Hz
Protons and electrons can be given very high energies in particle accelerators.

a. What is the wavelength in meters of an electron (mass = 9.11×10^−31kg) that has been accelerated to 99% of the speed of light?

b. In what region of the electromagnetic spectrum is this wavelength?
a. λ =
2.45×10−12
m

b. γ ray
What is the de Broglie wavelength (in meters) of a baseball weighing 145 g and traveling at 162 km/h ?
λ =
1.02×10−34
m
At what speed (in meters per second) must a 145 g baseball be traveling to have a de Broglie wavelength of 0.500 nm ?
v =
9.14×10−24
m/s
Use the Heisenberg uncertainty principle to calculate the uncertainty (in meters) in the position of a honeybee weighing 0.68 g and traveling at a velocity of 0.90 m/s . Assume that the uncertainty in the velocity is 0.1 m/s.
Δx≥
8×10−31
m
What are the four quantum numbers?

What does each specify?
the magnetic quantum number

the principal quantum number

the angular-momentum quantum number

the spin quantum number
What is meant by the term effective nuclear charge, Zeff?

What is it due to?
The effective nuclear charge is the net charge actually felt by an electron.

The electrons are shielded from the nucleus by the other electrons.
Give the allowable combinations of quantum numbers for each of the following electrons: A 3p electron
n = 3; l = 1; ml = -1, 0, +1; ms = ±½
Give the allowable combinations of quantum numbers for each of the following electrons: A 2p electron
n = 2; l = 1; ml = -1, 0, +1; ms = ±½
Give the allowable combinations of quantum numbers for each of the following electrons: A 3d electron
n = 3; l = 2; ml = -2, -1, 0, +1, +2; ms = ±½
Give the allowable combinations of quantum numbers for each of the following electrons: A 4p electron
n = 4; l = 1; ml = -1, 0, +1; ms = ±½
The wavelength of light at which the Balmer series converges corresponds to the amount of energy required to completely remove an electron from the second shell of a hydrogen atom.

Calculate this energy in kilojoules per mole.
E =
328.1
kJ/mol
Sodium atoms emit light with a wavelength of 330 nm when an electron moves from a 4p orbital to a 3s orbital.

What is the energy difference between the orbitals in kilojoules per mole?
E =
363
kJ/mol
Which orbital in each of the following pairs is higher in energy?

a. 5p or 5d

b. 4s or 3p

c. 6s or 4d
a. 5p

b. 4s

c. 6s
According to the aufbau principle, which orbital is filled immediately after each of the following in a multielectron atom?

a. 4s

b. 3d

c. 5f

d. 5p
a. 3d

b. 4p

c. 6d

d. 6s
Give the expected ground-state electron configurations for the following elements. Ti
Give the expected ground-state electron configurations for the following elements. Ru
Give the expected ground-state electron configurations for the following elements. Sn
Give the expected ground-state electron configurations for the following elements. Sr
Give the expected ground-state electron configurations for the following elements. Se
Give the expected ground-state electron configurations for the following elements:

a. Rb

b. W

c. Ge

d. Zr
a. [Kr]5s^1

b. [Xe]6s^24f^145d^4

c. [Ar]4s^23d^104p^2

d. [Kr]5s^24d^2
Order the electrons in the following orbitals according to their shielding ability: 4s, 4d, 4f.
4s>4d>4f
How many unpaired electrons are present in each of the following ground-state atoms?

a. O

b. Si

c. K

d. As
a. 2

b. 2

c. 1

d. 3
At what atomic number is the filling of a g orbital likely to begin?
Z =
121
Orbital energies in single-electron atoms or ions, such as He+, can be described with an equation similar to the Balmer-Rydberg equation:
(see picture)
where Z is the atomic number.

What wavelength of light in nm is emitted when the electron in He+ falls from n = 3 to n = 2?
λ =
164 nm
Give the expected ground-state electron configurations for the Sr.
Give the expected ground-state electron configurations for the Cd.
Give the expected ground-state electron configurations for atom with the following atomic number Z = 22.
Give the expected ground-state electron configurations for atom with the following atomic number Z = 34.
Subatomic Particles
Atoms are composed of three fundamental particles. Protons are positively charged, neutrons are neutral, and electrons are negatively charged. Protons and neutrons are clustered into a dense core called the nucleus, whereas electrons are found outside of the nucleus at a relatively large distance. Elements differ from one another by how many protons their atoms they contain. The number of protons is called the atomic number (Z) of the element. Since protons and neutrons make up most of the mass of an atom, the sum of the protons and neutrons is its mass number (A). In neutral atoms, the numbers of protons and electrons are equal. In ions, the numbers of electrons and protons are not equal.
Specify the number of protons, neutrons, and electrons in the neutral atom chromium-52.
24,28,24 protons, neutrons, electrons

All chromium atoms have 24 protons. A neutral Cr atom also has 24 electrons. The sum of the protons and neutrons gives the mass number,24+28=52.
The ion N3− has _____ protons and _____ electrons.
7,10

Regardless of how many electrons are present, every nitrogen nucleus contains 7 protons. The identity of an element is determined by the number of protons, not the number of electrons. When the element is neutral, the number of positively charged protons and negatively charged electrons will be equal. When there are more electrons than protons, the ion will be negative, as in this example. When there are more protons than electrons, the ion will be positive.
What isotope has 14 protons and 14 neutrons?
Silicon has 14 protons in its nucleus. It is the only element that has 14 protons. If it had more or fewer protons, it would not be silicon. However, the number of neutrons can vary, which is why the mass is written in the name.
Which element does X represent in the following expression: ³²₁₅X?
In this example we only needed the number of protons to identify the symbol of the element. If you were asked to name the isotope, the mass would be needed as well. In that case the correct answer would be phosphorus-32.
Ions and the Periodic Table
A main group metal tends to lose electrons, forming a cation with the same number of electrons as the nearest noble gas in the periodic table. A main group nonmetal tends to gain electrons, forming an anion with the same number of electrons as the nearest noble gas. The various groups gain or lose electrons as summarized in the following table:
If the following elements were to form ions, they would attain the same number of electrons as which noble gas?
The Be²⁺ ion has 2 electrons, just like He. The ions Mg²⁺ and F⁻ each have 10 electrons like Ne. The S²⁻ and Ca²⁺ each have 18 electrons like Ar. The ions Br⁻ and Sr²⁺ each have 36 electrons like Kr.
If the following elements were to form ions, what is the expectant charge?
A certain element forms an ion with 54 electrons and a charge of +2. Identify the element.
Ba
Naming Ionic Compounds
Because the common names of many chemical compounds are not helpful in indicating the chemical nature of the compound (e.g., the common name "table salt" gives no clue that this compound is composed of sodium and chlorine, NaCl), systematic names have been introduced to accurately identify compounds. The naming of simple ionic compounds is relatively straightforward, once you understand the naming system.
Ionic compounds are named according to their cation first, followed by their anion. Once the cation is named, the anion is then named. The naming conventions are listed in the tables to the right.
Finally, although ionic compounds need to have an overall neutral charge, the number of cations or anions are not mentioned in the formula name. For example, AlCl₃ is aluminum chloride, not aluminum trichloride.
Cations
Anions
What is the systematic name of Mg(NO₃)₂?
magnesium nitrate
What is the systematic name of NH₄ClO₃?
ammonium chlorate
What is the systematic name of PbO?
Ionic Compound Nomenclature and Formulas
Ionic bonds form when one atom completely transfers one or more electrons to another atom, resulting in the formation of ions. Positively charged ions (cations) are strongly attracted to negatively charged ions (anions) by electrical forces. All chemical compounds can be named systematically by following a series of rules. Binary ionic compounds are named by identifying first the positive ion and then the negative ion. Naming compounds with polyatomic ions involves memorizing the names and formulas of the most common ones.
Give the systematic name for the compound Ba(NO₃)₂.
barium nitrate

Barium is a metal that has only one oxidation state. Therefore it is not necessary to write II in parentheses in the systematic name. Only those metals with more than one oxidation state must have its state written in the name.
Give the systematic name for the compound Fe₂(SO₄)₃.
ferric sulphate

Although it seems like a small difference, iron(II) and iron(III) behave much differently chemically. They even form different-colored compounds because of the number of electrons they have to bond. Therefore it is very important to specify which oxidation state is being used.
Enter the formula for the compound magnesium oxide.
MgO

All of the elements in group 2 form ions with a +2 charge. That is because all of these elements need to lose two electrons to gain stability. After losing the electrons, there is an excess of two protons, which results in a +2 charge.
Enter the formula for the compound lead(II) phosphate.
Pb₃(PO₄)₂

Lead, like iron, has more than one oxidation state, so it is necessary to put the oxidation state of the metal in parentheses. Elements in group 1 and 2 only have one oxidation state and so there is no need for Roman numerals in their systematic name.
Ionic Compound Formulas
Many chemical compounds have both common and systematic names. Common names are historical and tend not to identify the elements that make up the compound. However, the systematic name allows for correct identification of the cations and anions that together make up the ionic compound.
What is the chemical formula for potassium permanganate?
KMnO₄
Blue vitriol is commonly used in industrial dyeing processes. What is the chemical formula for blue vitriol, whose systematic name is copper(II) sulfate?
CuSO₄
Sodium carbonate is used in the manufacture of paper. What is the chemical formula for this compound?
Na₂CO₃
What is the correct ionic formula when Mg²⁺ and P³⁻ react?
Mg₃P₂
What is the correct ionic formula when Al³⁺ and SO₄²⁻ react?
Al₂(SO₄)₃
Three binary ionic compounds are represented on the following periodic table: red with red, green with green, and blue with blue (see picture).

a. What is likely formula of red ionic compound?

b. Name red ionic compound.

c. What is likely formula of green ionic compound?

d. Name green ionic compound.

e. What is likely formula of blue ionic compound?

f. Name blue ionic compound.
a. K₂S

b. potassium sulfide

c. SrI₂

d. strontium iodide

e. Ga₂O₃

f. gallium oxide
The following drawings are those of solid ionic compounds, with red spheres representing the cations and blue spheres representing the anions in each. Which of the following formulas are consistent with each drawing? (see picture)

a. drawing 1

b. drawing 2
a. CaCl₂

b. LiBr
NaNO₂
Give systematic names for the following compounds.

a. CsF

b. KBr

c. CuF₂

d. CuS

e. CuBr₂
a. cesium fluoride

b. potassium bromide

c. copper(II) fluoride

d. copper(II) sulfide

e. copper(II) bromide
Write formulas for the following compounds.

b. Manganese(II) chloride

c. Copper(II) oxide

d. Aluminum oxide
a. VCl₃

b. MnCl₂

c. CuO

d. Al₂O₃
Give systematic names for the following compounds.

a. LiCN

b. Ag₂S₂O₃

c. NaH₂PO₄

d. (Pb(ClO₄)₂

e. Sn(H₂PO₄)₄

f. (NH₄)₂SO₄
a. lithium cyanide

b. silver thiosulphate

c. sodium dihydrogen phosphate

e. tin(IV) dihydrogen phosphate

f. ammonium sulfate
Write formulas for the following binary compounds.

a. Potassium chloride

b. Tin(II) bromide

c. Calcium oxide

d. Barium chloride

e. Aluminum hydride
a. KCl

b. SnBr₂

c. CaO

d. BaCl₂

e. AlH₃
Name the following ions.

a. Ca²⁺

b. Cs⁺

c. Na⁺

d. HCO₃⁻

e. Hg⁺

f. Fe³⁺

g. CH₃CO₂⁻

h. Cr₂O₇²⁻

i. Mn²⁺

j. ClO₄⁻
a. calcium ion

b. cesium ion

c. sodium ion

d. hydrogen carbonate ion

e. mercury(I) ion

f. iron(III) ion

g. acetate ion

h. dichromate ion

i. manganese(II) ion

j. perchlorate ion
Mg²⁺ and Cl⁻
MgCl₂
Ca²⁺ and SO ²⁻₄
CaSO₄
Ga³⁺ and SO²⁻₄
Ga₂(SO₄)₃
Chlorine and magnesium
MgCl₂
Oxygen and magnesium
MgO
Sulfer and magnesium
MgS
Sulfate ion
SO₄²⁻
Phosphate ion
PO₄³⁻
Zirconium(IV) ion
Zr⁴⁺
Chromate ion
CrO₄²⁻
Acetate ion
C₂H₃O₂⁻
Thiosulfate ion
S₂O₃²⁻
Na?SO₄
Na₂SO₄
Ba?(PO₄)?
Ba₃(PO₄)₂
Ga?(SO₄)?
Ga₂(SO₄)₃
Electron Configurations of Ions
When an atom forms an ion, it will gain or lose electrons to attain a more stable electron configuration, frequently that of a noble gas. Nonmetals tend to form anions by gaining electrons, which enter the lowest energy unoccupied orbital. Metals tend to form cations by losing electrons. Main group metals lose electrons in the reverse order of filling. Transition metals, however, lose s electrons first.
In the ground-state electron configuration of Fe³⁺, how many unpaired electrons are present?
5
Build the orbital diagram for the ion most likely formed by phosphorus.
Electron Configurations of Atoms and Ions
The electron configuration of an atom tells us how many electrons are in each orbital. For example, helium has two electrons in the 1s orbital. Therefore the electron configuration of He is 1s².
What is the ground-state electron configuration of a neutral atom of cobalt?
[Ar]3d^74s^2
What is the ground-state electron configuration of the fluoride ion F⁻?
[He]2s^22p^6
Which element has the following configuration: [Kr]5s²?
Strontium
What doubly positive ion has the following ground-state electron configuration? 1s²2s²2p⁶3s²3p⁶3d⁶
Fe²⁺
First Ionization Energy
Ionization energy is the energy required to remove an electron from an atom or ion. First ionization energy refers to the energy required to remove an electron from an electrically neutral atom in the gas phase. Subsequent ionization energies reflect the energies needed to strip successive electrons off an increasingly positively charged ion.
Arrange the elements in decreasing order of first ionization energy.
Arrange the elements in order of decreasing first ionization energy.

element x 116 pm

element y 196 pm

element z 260 pm
The inward "pull" on the electrons from the nucleus is called the effective nuclear charge. The greater the pull on the electrons, the smaller the molecule and the harder it is to remove an electron.
Which of the following arrangements of the elements Cl, F, S, Sn, and Te is in order of decreasing ionization energy?
F, Cl, S, Te, and Sn
Using the periodic table as your guide, predict which element in each of the following pairs has the larger ionization energy.

a. K or Br

b. S or Te

c. Ga or Se

d. Ne or Sr
a. Br

b. S

c. Se

d. Ne
Na?SO₄
Na₂SO₄
Ba?(PO₄)?
Ba₃(PO₄)₃
Ga?(SO₄)?
Ga₂(SO₄)₃
a. Order the indicated three elements according to the ease with which each is likely to lose its third electron:
Most ease: Al, Kr, Ca: least ease
Three atoms have the following electron configurations:
1s^22s^22p^63s^23p¹

1s^22s^22p^63s^23p^5

1s^22s^22p^63s^23p^64s¹

a. Which of the three has the largest Ei1?

b. Which has the smallest Ei4?
a. 1s^22s^22p^63s^23p^5

b. 1s^22s^22p^63s^23p^5
Electron affinity is the measure of the attraction of an electron toward an isolated gaseous atom. When an electron is added to an isolated gaseous atom or ion energy is either released or absorbed; this energy change is known as electron affinity. Electron affinity is positive when energy is absorbed, and it is negative when energy is released.
Electron affinity is the energy change that occurs when an electron is added to a gaseous atom or ion. The electron affinity of an O atom is −142 kJ, meaning that when an electron is added to an O atom, energy is released and the O− ion is stable. However, when a second electron is added to an O− ion, energy is required. Since the O− ion is already negatively charged, adding another negatively charged electron is difficult. Thus, the electron affinity of an O− ion is +710 kJ.
Electron affinity and electron configuration of the atoms
The electron affinity of elements is a periodic property, so you can predict whether the electron affinity of an element is positive or negative based on its electron configuration.
In general, you can use the following rules to correlate electron affinity and electron configuration: (see picture)
The electron affinity of an element is positive if energy is absorbed during the process of accepting an electron and negative if energy is released during the process. Using the electron configuration of the atom or ion undergoing the addition of an electron, predict whether the electron affinity will be positive or negative for the following reactions, and classify them accordingly.
Consider the addition of an electron to the following atoms from the third period. Rank the atoms in order from the most negative to the least negative electron affinity values based on their electron configurations.
The electron affinity values for Cl and Si are −349 kJ/mol and −134 kJ/mol, respectively. The noble gas (group 8A) has a positive electron affinity. This suggests that Cl has the most negative electron affinity value, whereas Ar has the least negative electron affinity value.
Electron Affinity and Electron Configurations
Electron affinity, Eea, is the change in energy that occurs when an electron is added to a neutral isolated atom. This can be represented by the following equation:
X(g)+e−→X−(g)
Most electron affinity values are negative because energy is usually released when a neutral atom gains an electron. Eea values become more negative with increasing tendency of the atom to accept an electron and increasing stability of the resulting anion. Eea shows a periodic trend that is related to electron configuration. Elements with less than an octet and with high effective nuclear charge (Zeff) tend to have large negative Eea values. Elements with filled valence shells or subshells and low Zeff tend to have Eea values near zero.
Consider the following neutral electron configurations in which n has a constant value. Which configuration would belong to the element with the most negative electron affinity, Eea?
2s22p5
Arrange the following elements from greatest to least tendency to accept an electron.
The tendency to gain an electron is quantitatively measured by the electron affinity, the amount of energy involved in the addition of an electron to a neutral gaseous atom. Ordering these elements by the electron affinity provides an identical order:
F>O>C>Li>Be
Periodic Trends in Relative Electron Affinity
Electron affinity, EA, is the energy required to add an electron to a neutral gaseous atom and is related to an element's position on the periodic table. Electron affinities can be positive, negative, or zero, as shown in the table.
For the elements with the electron affinities given in the table in the introduction, which element is most likely to accept an electron?
Br

Because bromine is only one electron away from noble-gas configuration, it becomes incredibly stable upon the addition of that one electron. Even though nitrogen also needs to gain electrons to achieve a noble-gas configuration, N− is nowhere near as stable as Br−. Calcium must lose electrons to achieve noble-gas configuration. Therefore, adding an electron to a neutral calcium atom makes it less stable.
Rank the following elements by electron affinity, from most positive to most negative EA value.
The reason why group 8A elements have a positive EA value is because they are incredibly stable in their neutral form with an octet of electrons in the outermost energy level. They have little tendency to gain an electron. The reason that group 5A elements have an EA value that is less negative than expected is because a half-filled p subshell is particularly stable.
Which of the following electron configurations corresponds to an element with the most positive electron affinity?
1s 22s 22p 63s 23p 6
Which of the indicated three elements has the least favorable Eea?
Kr
Which of the indicated three elements has the most favorable Eea?
Ge
The size of ions as measured by ionic radii varies in a systematic manner. The size of the ion can be explained in part by effective nuclear charge, Zeff, which is the net nuclear charge felt by an electron. The effective nuclear charge takes into account the actual nuclear charge and the shielding of this charge by inner electrons. When an atom loses electrons, the resulting cation is smaller both because the remaining electrons experience a larger Zeff and because these electrons are usually in orbitals closer to the nucleus than the electrons that were lost. The more electrons that are lost, the smaller the ion becomes.
Similarly, when an atom gains electrons, the resulting anion is larger owing to both increased electron-electron repulsions and a reduction in Zeff. The more electrons that are gained, the larger the ion becomes.
Rank the following ions in order of decreasing radius: Rb+,K+,Li+,Cs+, and Na+. Use the periodic table as necessary.
Rank the following items in order of decreasing radius: Mg, Mg2+, and Mg2−.
The following ions contain the same number of electrons. Rank them in order of decreasing ionic radii.
In a group of ions with the same number of electrons, the most negative ion is the largest and the most positive ion is the smallest.
Which of the following spheres represents a K+ ion, which a K atom, and which a Cl− ion?
Ionic bonds generally form between ____.
an element with a small E i and an element with a large negative E ea
Which substance in each of the following pairs has the larger lattice energy?

KCl or RbCl
KCl
Which substance in each of the following pairs has the larger lattice energy?

CaF2 or BaF2
CaF₂
Which substance in each of the following pairs has the larger lattice energy?

CaO or KI
CaO
One of the following pictures(Figure 1) represents NaCl and one represents MgO

a. Which is which?

b. Which has the larger lattice energy?
a. a) is NaCl and (b) is MgO

b. MgO
What are group 6A elements likely to do when they form ions-gain electrons or lose them?

How many?
gain electrons

2
What noble gas configurations are the following elements likely to adopt in reactions when they form ions?

a. Rb

b. Ba

c. Ga

d. F
a. Kr

b. Xe

c. Ar

d. Ne
Which of the following factors has no effect on an element's gaining or losing electrons to form an octet?
A low Z eff makes it easy to add electrons to form anions with eight valence-shell electrons.
In the following drawings, red spheres represent cations and blue spheres represent anions. Match each of the drawings (Figure 1) - with the following ionic compounds.

a. Ca₃(PO₄)₂

b. Li₂CO₃

c. FeCl₂

d. MgSO₄
a. d

b. b

c. c

d. a
Which of the following drawings(Figure 1) is more likely to represent an ionic compound, and which a covalent compound?
(a) is an ionic compound, (b) is a covalent compound
Circle the approximate part or parts of the periodic table where the following elements appear.

(a) Elements with the smallest values of Ei1

(b) Elements with the largest atomic radii

(c) Elements with the most negative values of Eea
Where on the periodic table would you find the element that has an ion with each of the following electron configurations? Identify each ion.

(a) 3+ ion: 1s22s22p6

(b) 3+ ion: [Ar]3d3

(c) 2+ ion: [Kr]5s24d10

(d) 1+ ion: [Kr]4d10
Which of the following spheres is likely to represent a metal, and which a nonmetal?

Which sphere in the products represents a cation, and which an anion?
Green sphere represents a nonmetal, blue sphere represents a metal.

Green sphere represents an anion, blue sphere represents a cation.
Each of the pictures (a)-(d) represents one of the following substances at 25 ∘C: sodium, chlorine, iodine, sodium chloride.

Which picture corresponds to which substance?
a. Picture (a) corresponds to: Iodine

b. Picture (b) corresponds to: Sodium

c. Picture (c) corresponds to: Sodium chloride

d. Picture (d) corresponds to: Chlorine
Three binary compounds are represented on the following drawing: red with red, blue with blue, and green with green.

a. Give a likely formula for red compound.

b. Give a likely formula for blue compound.

c. Give a likely formula for green compound.
a. PbS₂

b. SrF₂

c. CBr₄
What is the difference between a covalent bond and an ionic bond?
A covalent bond results when two atoms share several (usually two) of their electrons.

An ionic bond results from a complete transfer of one or more electrons from one atom to another.
Give an example of an ionic bond.
The bond in LiF (Li+F−)

The bond in NaCl (Na+Cl−)
Give an example of a covalent bond.
The Si−H bonds in SiH4

The B−H bonds in BH3
How many protons and electrons are in each of the following ions?

a. Se²⁻

b. Au³⁺

c. Be²⁺

d. Rb⁺
a. 34p, 36e

b. 79p, 76e

c. 4p, 2e

d. 37p, 36e
What is the identity of the element X in the following ions?

a. X²⁺, a cation that has 36 electrons.

b. X⁻, an anion that has 36 electrons.
a. Sr

b. Br
Write formulas for the following compounds.

a. Calcium phosphate

b. Barium hydrogen sulfate

c. Manganese(II) nitrate

d. Chromium(III) phosphate
a. Ca₃(PO₄)₂

b. Ba(HSO₄)₂

c. Mn(NO₃)₂

d. CrPO₄
Write formulas for the following compounds.

a. Calcium acetate

b. Iron(II) cyanide

c. Calcium chlorite

d. Barium perchlorate

e. Aluminum sulfite
a. Ca(C₂H₃O₂)₂

b. Fe(CN)₂

c. Ca(ClO₂)₂

d. Ba(ClO₄)₂

e. Al₂(SO₃)₃
Name each of the following compounds.

a. MgSO₄

b. ZnCrO₄

c. Na₂CO₃

d. LiClO₄

e. Ca₃(PO₄)₂

f. KMnO₄
a. magnesium sulfate

b. zinc chromate

c. sodium carbonate

d. lithium perchlorate

e. calcium phosphate

f. potassium permanganate
What are the formulas of the compounds formed from the following ions:

a. Ca²⁺ and Br⁻

b. K⁺ and SO₄²⁻

c. Al³⁻ and SO₄²⁻
a. CaBr₂

b. K₂SO₄

c. Al₂(SO₄)₃
Write formulas for compounds of rubidium with each of the following elements:

a. Bromine

b. Nitrogen

c. Sulfer
a. RbBr

b. Rb₃N

c. Rb₂S
What are the charges on the positive ions in the following compounds?

a. Zn(CN)₂

b. Fe(NO₂)₃

c. Ti(SO₄)₂

d. Sn₃(PO₄)₂

e. Hg₂S

f. MnO₂

g. KIO₄

h. Cu(CH₃CO₂)₂
a. Zn²⁺

b. Fe²⁺

c. Ti⁴⁺

d. Sn²⁺

e. Hg₂²⁺

f. Mn⁴⁺

g. K⁺

h. Cu²⁺
Write formulas for each of the following compounds.

a. Aluminum bromide

b. Chromium(III) sulfate

c. Sodium peroxide
a. AlBr₃

b. Cr₂(SO₄)₃

c. Na₂O₂
What are the likely ground-state electron configurations of the following cations?

a. La³⁺

b. Ag⁺

c. Sn²⁺
a. [Xe]

b. [Kr]4d^10

c. [Kr]5s^24d^10
What is the electron configuration of Ca²⁺?
What is the electron configuration of Ti²⁺?
Identify the element whose 2+ ion has the ground-state electron configuration [Ar] 3d¹⁰.
Zn
There are two elements in the transition metal series Sc through Zn that have four unpaired electrons in their 2+ ions.

Identify the elements.
Cr, Fe
What tripositive ion has the electron configuration [Kr]4d³?
Mo³⁺
What neutral atom has the electron configuration [Kr]5s²4d¹?
Y
Which group of elements in the periodic table has the largest Ei1?
Group 8A
Which group of elements in the periodic table has the smallest Ei1?
Group 1A
Which element in the periodic table has the smallest ionization energy?
Cs
Which element in the periodic table has the largest ionization energy?
He
Which has the smaller second ionization energy, Rb or Sr?
Sr
Which has the larger third ionization energy, In or Sr?
Sr
Which has the smaller fourth ionization energy, Sn or Sb?
Sn
Which has the larger sixth ionization energy, Se or Br?
Br
Three atoms have the following electron configurations.

1s22s22p63s23p3

1s22s22p63s23p6

1s22s22p63s23p64s2

a. Which of the three has the largest Ei2?

b. Which has the smallest Ei7?
a. 1s22s22p63s23p6

b. 1s22s22p63s23p6
Write the electron configuration of the atom in the third row of the periodic table that has the smallest Ei4.
The first four ionization energies in kJ/mol of a certain second-row element are 801, 2427, 3660, and 25,025.

What is the likely identity of the element?
B
The first four ionization energies in kJ/mol of a certain second-row element are 900, 1757, 14,849, and 21,007.

What is the likely identity of the element?
Be
Which element in each of the following sets has the smallest first ionization energy? Which has the largest?

a. Li, Ba, K

b. B, Be , Cl

c. Ca, C, Cl
a. K - lowest, Li - highest

b. B - lowest, Cl - highest

c. Ca - lowest, Cl - highest
What elements meet the following descriptions?

a. Has largest Ei3

b. Has largest Ei7
a. Group 2A

b. Group 6A
What is the relationship between the ionization energy of a monoanion such as Cl− and the electron affinity of the neutral atom?
This quantities have the same magnitude but opposite sign.
What is the relationship between the electron affinity of a singly charged cation such as Na+ and the ionization energy of the neutral atom?
The relationship between the electron affinity of a univalent cation and ionization energy of the neutral atom is that they have the same magnitude but opposite signs
Which has the more negative electron affinity, Br or Br−?
Br
Which has the more negative electron affinity, Na+ or Na?
Na⁺
Which has the more negative electron affinity, Na+ or Cl?
Na⁺
Why is energy usually released when an electron is added to a neutral atom but absorbed when an electron is removed from a neutral atom?
Energy is usually released when electron is added to a neutral atom but absorbed when an electron is removed from a neutral atom because of the positive Zeff.
Why does ionization energy increase regularly across the periodic table from group 1A to group 8A, whereas electron affinity increases irregularly from group 1A to group 7A and then falls dramatically for group 8A?
The electron affinity increases irregulary from 1A to 7A and then falls dramatically for Group 8A because the additional electron goes into the next higher shell.

Ei1 increases steadily across the periodic table from Group 1A to Group 8A because electrons are being removed from the same shell and Zeff is increasing.
Why does phosphorus have a less-negative electron affinity than its neighbors silicon and sulfur?
The 3p orbitals in P are half-filled. The electron affinity for Si is more negative because the added electron is going into an empty 3p orbital. The electron affinity for S is more negative because of a higher Zeff.
Order the following compounds according to their expected lattice energies: LiCl, KCl, KBr, MgCl₂.
MgCl₂,LiCl,KCl,KBr
Which element, indicated by letter on the periodic table above, has a 3+ ion with the electron configuration 1s 22s 22p 63s 23p 6?
A
The correct chemical formula for manganese(IV) acetate is
Mn(CH₃CO₂)₄
Which of the following represents the change in electronic configuration that is associated with the first ionization energy of magnesium?
[Ne]3s²→[Ne]3s¹+e-
Which of the following ionic compounds would be expected to have the highest lattice energy?

LiCl
LiBr
LiI
LiF
LiF
Which of the following is the correct chemical formula for a molecule of chlorine?
Cl₂
The four spheres below represent Na+, Mg2+, F-, and O2-, not necessarily in that order.

(see picture)

Which sphere most likely represents the F- ion?
C
Of the following, which element has the highest first ionization energy?

Se
Na
Ca
Cl
Cl

See link for all ionization energies

https://en.wikipedia.org/wiki/Ionization_energies_of_the_elements_(data_page)
How many electrons are in the outermost shell of the Sn4+ ion in its ground state?
18
What is the identity of element Q if the ion Q²⁺ contains 10 electrons?
Mg (it normally has 12 electrons).
Li₂S is named
lithium sulfide
The following four spheres represent a metal atom, a nonmetal atom, a monatomic anion and a monatomic cation, not necessarily in that order.

(see picture)

In the reaction shown, which sphere most likely represents the metal atom?
A

(B is a non metal atom, C is a cation, and D is an anion)
Which element has the most favorable (most negative) electron affinity?

O
Ne
Mg
Na
O
Which of the following represents the change in electronic configuration that is associated with the first ionization energy of barium?
[Xe]6 s 2 → [Xe]6 s 1 + e-
The ion NO₂⁻ is named
Nitrite ion
Which of these elements has the most favorable (most negative) electron affinity?

Ca
S
N
Ne
S
To reach a noble gas electron configuration how many electrons would sulfur have to adopt?
2
Atoms of which element, indicated by letter on the periodic table, is expected to have the smallest atomic radius?
C

Chart of atomic radii

Which element has the highest (most negative) electron affinity?

Xe
Cs
Ba
Te
Te
Of the following, which element has the highest first ionization energy?