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CHE 101: Atoms and Elements
Terms in this set (82)
3 elements for growth: potassium, nitrogen, and phosphorus.
All matter is composed of
elements, of which there are 118 different kinds. Of these, 88 elements occur naturally and make up all substances in our world.
In our bodies,
Calcium and phosphorus form the structure of bones and teeth, iron and copper are needed in the formation of red blood cells, and iodine is required for thyroid function.
Chemical symbols are
1 or 2 letter abbreviations for the names of the elements. Only the first letter of an element's symbol is capitalized. If the symbol has a second letter, it is lowercase so that we know when a different element is indicated. If 2 letters are capitalized, they represent the symbols of 2 different elements.
By the late 1800s,
scientists recognized that certain elements looked alike and behaved the same way.
In 1869, a Russian chemist,
Dmitri Mendeleev, arranged the 60 elements known at the time into groups with similar properties and placed them in order of increasing atomic masses.
Each vertical column on the periodic table contains a
group (or family) of elements that have similar properties.
A group number is written at the top of each vertical column (group) in the periodic table.
For many years, the
representative elements have had group numbers 1A to 8A.
In the center of the periodic table is a block of elements known as the
transition elements, which have numbers followed by the letter "B".
The 2 rows of 14 elements called the lanthanides and actinides (or the inner transition elements) are placed at the bottom of the periodic table to allow them to fit on a page.
Each horizontal row in the periodic table is a
period. The periods are counted down from the top of the table as Periods 1 to 7.
The first period contains 2 elements: hydrogen (H) and helium (He).
The second period contains 8 elements: Lithium (Li), beryllium (Be), boron (B), carbon (C), nitrogen (N), oxygen (O), fluorine (F), and neon (Ne).
The third period also contains 8 elements beginning with sodium (Na), and ending with argon (Ar). The fourth period which begins with potassium (K), and the fifth period, which begins with cesium (Cs), has 32 elements. The seventh period contains 32 elements, for a total of 118 elements.
Several groups in the periodic table have special names -
Group 1A - (1) elements - lithium, sodium, potassium, rubidium (Rb), cesium, and francium (Fr) - are a family of elements known as the alkali metals.
The elements within this group are soft, shiny metals that are good conductors of heat and electricity and have relatively low melting points. Alkali metal react vigorously with water and form white products when they combine with oxygen (O2).
Although hydrogen is at the top of group 1A (1), it is not an alkali metal and has very different properties than the rest of the elements in this group. Thus, hydrogen is not included in the alkali metals.
The alkaline earth metals are found in Group
2A (2). They include the elements beryllium, magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). The alkaline metals are shiny metals like those in group 1A (1), but they are not as reactive.
The halogens are found on the right side of the periodic table in group
7A (17). They include the elements fluorine, chlorine (Cl), bromine (Br), iodine (I), astatine (At), and tennessine (Ts). The halogens, especially fluorine and chlorine, are highly reactive and form compounds with most of the elements.
The noble gases are found in group
8A (18). They include helium, neon, argon (Ar), krypton (Kr), xenon (Xe), radon (Rn), and oganesson (Og). They are quite unreactive and are seldom found in combination with other elements.
Another feature of the periodic table is the heavy zigzag line that separates the elements into the
metals and the nonmetals. Except for hydrogen, the metals are the the left of the line with the nonmetals to the right.
In general, most metals are
shiny solids, such as copper (Cu), gold (Au), and silver (Ag). Metals can be shaped into wires (ductile) or hammered into a flat sheet (malleable). Metals are good conductors of heat and electricity. They usually melt at higher temperatures than nonmetals. All metals are solid at room temperature, except for mercury (Hg), which is a liquid.
Nonmetals are not especially
shiny, ductile, or malleable, and they are often poor conductors of heat and electricity. They typically have low melting points and low densities. Some examples of nonmetals are hydrogen, carbon, nitrogen, oxygen, chlorine, and sulfur (S).
Except for aluminum and oganesson, the elements located along the heavy line are
metalloids: B, Si, Ge, As, Sb, Te, Po, At, and Ts.
Metalloids are elements that exhibit some properties that are typical of the metals and other properties that are characteristic of the nonmetals. For example, they are better conductors of heat and electricity than the nonmetals, but not as good as the metals. The metalloids are semiconductors because they can be modified to function as conductors or insulators.
Of all the elements, only about 20 are essential for the well-being and survival of the human body
Of those, 4 elements, oxygen, carbon, hydrogen, and nitrogen - which are representative elements in period 1 and period 2 on the periodic table, make up 96% of our body mass. Most of the food in our daily diet provides these elements to maintain a healthy body. These elements are found in carbs, fats, and proteins. Most of the hydrogen and oxygen is found in water, which makes up 55-60% of our body mass.
The macrominerals -
Ca, P, K, Cl, S, Na, and Mg - are located in Period 3 and Period 4 of the periodic table. They are involved in the formation of bones and teeth, maintenance of heart and blood vessels, muscle contraction, nerve impulses, acid-base balance of body fluids, and regulations of cellular metabolism. The macrominerals are present in lower amounts than the major elements, so that smaller amounts are required in our daily diets.
The other essential elements, called microminerals or trace minerals, are mostly
transition elements in Period 4 along with Si in Period 3 and Mo and I in Period 5. They are present in the human body in very small amounts, some less than 100 mg. In recent years, the detection of such small amounts has improved so that researchers can more easily identify the roles of trace elements. Some trace elements such as arsenic, chromium, and selenium are toxic at high levels in the body but are still required by the body. Other elements, such as tin and nickel, are thought to be essential, but their metabolic role has not yet been determined.
An atom is the
smallest particle of an element that retains the characteristics of that element.
Greek philosophers in 500 B.C.E. reasoned that
everything must contain minute particles called atomos, the idea of atoms did not become a scientific theory until 1808. Then, John Dalton (1766-1844) developed an atomic theory that proposed that atoms were responsible for the combinations of elements found in compounds.
Dalton's Atomic Theory:
1. All matter is made up of tiny particles called atoms.
2. All atoms of a given element are the same and different from atoms of other elements.
3. Atoms of 2 or more different elements combine to form compounds. A particular compound is always made up of the same kinds of atoms and always has the same number of each kind of atom.
4. A chemical reaction involves the rearrangement, separation, or combination of atoms. Atoms are neither created nor destroyed during a chemical reaction.
Dalton's atomic theory formed the basis of current
atomic theory, although we have modified some of Dalton's statements. We now know that atoms of the same elements are not completely identical to each other and consist of even smaller particles. However, an atom is still the smallest particle that retains the properties of an element.
A special kind of microscope called a
scanning tunneling microscope (STM) produces images of individual atoms.
By the end of the 1800s, experiments showed that atoms were not solid spheres but were composed of smaller bits of matter called
subatomic particles, 3 of which are the proton, neutron, and electron. 2 of these subatomic particles were discovered because they have electrical charges.
In 1897, J.J. Thomson, an English physicist, applied electricity to electrodes sealed in a glass tube, which produced streams of small particles called
cathode rays. Because these rays were attached to a positively charged electrode, Thomson realized that the particles in the rays must be negatively charged. In further experiments, these particles called electrons were found to be much smaller than the atom and to have extremely small masses. Because atoms are neutral, scientists soon discovered that atoms contained positively charged particles called protons that were much heavier than the electrons.
Thomson proposed a "plum-pudding" model for the atom in which the electrons and protons were randomly distributed in a positively charged cloud like plums in a pudding.
In 1911, Ernest Rutherford worked with Thomson to test this model. In Rutherford's experiment, positively charged particles were aimed at a thin sheet of gold foil. If the Thomson model were correct, the particles would travel in straight paths through the gold foil. Rutherford was greatly surprised to find that some of the particles were deflected as they passed through the gold foil, and a few particles were deflected so much that they went back in the opposite direction. According to Rutherford, it was as though he had shot a cannonball at a piece of tissue paper, and it bounced back at him.
From his gold-foil experiments, Rutherford realized that the
protons must be contained in a small, positively charged region at the center of the atom, which he called the nucleus. He proposed that the electrons in the atom occupy the space surrounding the nucleus through which most of the particles traveled undisturbed. Only the particles that came near this dense, positive center were deflected. If an atom were the size of a football stadium, the nucleus would be about the size of a golf ball placed in the center of the field.
In an atom,
the protons and neutrons that make up almost all the mass are packed into the tiny volume of the nucleus. The rapidly moving electrons (negative charge) surround the nucleus and account for the large volume of the atom.
One proton has a mass of 1.67 x 10^-24 g, and the neutron is about the same. However, the electron has a mass of 9.11 x 10^-28 g, which is much less than the mass of either a proton or neutron. Because the masses of subatomic particles are so small, chemists use a very small unit of mass called an
atomic mass unit (amu). An amu is defined as one-twelfth of the mass of a carbon atom, which has a nucleus containing 6 protons and 6 neutrons. In biology, the atomic mass unit called a Dalton (Da) in the honor of John Dalton. On the amu scale, the proton and neutron each have a mass of about 1 amu. Because the electron mass is so small, it is usually ignored in atomic mass calculations.
All the atoms of the same element always have the same number of protons.
This features distinguishes atoms of 1 element from atoms of all the other elements.
The atomic number of an element is equal to the
number of protons in every atom of that element. The atomic number is the whole number that appears above the symbol of each element on the periodic table.
Atomic # = number of protons in an atom.
We can use an atomic number to identify the number of protons in an atom of any element. For example, a lithium atom, with atomic number 3, has 3 protons.
Any atom with 3 protons is always a lithium atom. In the same way, we determine that a carbon atom, with atomic number 6, has 6 protons. Any atom with 6 protons is carbon.
An atom is electrically neutral. That means that the number of protons in an atom is equal to the number of
electrons, which gives every atom an overall charge of zero. Thus, the atomic number also gives the number of electrons.
We now know that the protons and neutrons determine the mass of the nucleus. Thus, for a single atom, we assign a
mass number, which is the total number of protons and neutrons in its nucleus. However, the mass number does not appear on the periodic table because it applies to single atoms only.
Mass number = number of protons + number of neutrons
Number of neutrons in a nucleus = mass number - number of protons
Carbon has the symbol C. However, its atom can be rearranged different ways to give several different substances. 2 forms of carbon - diamond and graphite - have been known since prehistoric times.
A diamond is transparent and harder than any other substance, whereas graphite is black and soft. In diamond, carbon atoms are arranged in a rigid structure. In graphite, carbon atoms are arranged in flat sheets that slide over each other. Graphite is used as a pencil lead and as a lubricant.
2 other forms of carbon have been discovered more recently in the formed called Buckminsterfullerene or buckyball (named after R. Buckminster Fuller, who popularized the geodesic dome), 60 carbon atoms are arranged as rings of 5 and 6 atoms to five a spherical, cage-like structure.
When a fullerene structure is stretched out, it produces a cylinder with a diameter of only a few nanometers called a nanotube. Partical uses for buckyballs and nanotubes are not yet developed, but they are expected to find use in lightweight structural materials, heat conductors, computer parts, and medicine. Recent research has shown that carbon nanotubes (CNT) can carry many drug molecules that can be released once the CNT enter the targeted cells.
However, the atoms of any one element are not identical because the atoms of most elements have different numbers of neutrons.
When a sample of an element consists of 2 or more atoms with differing numbers of neutrons, those atoms are called isotopes.
Isotopes are atoms of the same element that have the same atomic number but different numbers of neutrons.
For example, all atoms of the element magnesium (Mg) have an atomic number of 12. Thus every magnesium atom always has 12 protons. However, some naturally occurring magnesium atoms have 12 neutrons, others have 13 neutrons, and still others have 14 neutrons. The different numbers of neutrons give the magnesium atoms different mass numbers but do not change their chemical behavior. The 3 isotopes of magnesium have the same atomic number but different mass numbers.
To distinguish between the different isotopes of an element, we write an atomic symbol for a particular isotope that indicates the mass number in the upper left corner and the atomic number in the lower left corner.
An isotope may be referred to by its name or symbol, followed by its mass number, such as magnesium-24 or Mg-24. Magnesium has 3 naturally occurring isotopes. In a large sample of naturally occurring magnesium atoms, each type of isotope can be present as a low percentage or a high percentage. For example, the Mg-24 isotope makes up almost 80% of the total sample, whereas Mg-25 and Mg-26 each make up only about 10% of the total number of magnesium atoms.
Because each isotope has a different mass, chemists have calculated an atomic mass for an "average atom",
which is a weighted average of the masses of all the naturally occurring isotopes of that element. On the periodic table, the atomic mass is the number including decimal places that is given below the symbol of each element. Most elements consist of 2 or more isotopes, which is one reason that the atomic masses on the periodic table are seldom whole numbers.
To understand how the atomic mass as a weighted average for a group of isotopes is calculated, we will use an analogy of bowling balls with different weights.
To calculate the atomic mass of an element, we need to know the percentage abundance and the mass of each isotope, which must be determined experimentally. For example, a large sample of naturally occurring chlorine atoms consists of 75.76% of 35/17 Cl atoms and 24.24% of 37/17 Cl atoms. The 35/17 Cl isotope has a mass of 34.97 amu, and the 37/17 Cl isotope has a mass of 36.97 amu.
When we listen to a radio, use a microwave oven, turn on a light, see the colors of a rainbow, or have an X-ray taken, we are experiencing various forms of
In an electromagnetic wave, just like the waves in an ocean, the distance between the peaks is called the wavelength.
All forms of electromagnetic radiation travel in space at the speed of light, 3.0 x 10^8 m/s, but differ in energy and wavelength. High-energy radiation has short wavelengths compared to low-energy radiation, which has longer wavelengths. The electromagnetic spectrum shows the arrangement of different types of electromagnetic radiation in order of increasing energy.
The light energy, especially ultraviolet (UV), excites electrons and may lead to unwanted chemical reactions. The list of damaging effects of sunlight includes sunburn; wrinkling; premature aging of the skin; changes in the DNA of the cells, which can lead to skin cancers; inflammations of the eyes; and perhaps cataracts.
Some drugs, like the acne medications Accutane and Retin-A, as well as antibiotics, diuretics, sulfonamides, and estrogen, make the skin extremely sensitive to light.
Phototherapy uses light to treat certain
skin conditions, including psoriasis, eczema, and dermatitis. In the treatment of psoriasis, for example, oral drugs are given to make the skin more photosensitive; then exposure to UV radiation follows. Low-energy radiation (blue light) with wavelengths from 390 to 470 nm is used to treat babies with neonatal jaundice, which converts high levels of bilirubin to water-soluble compounds that can be excreted from the body. Sunlight is also a factor in stimulating the immune system.
In a disorder called seasonal affective disorder, or SAD, people experience
mood swings and depression during the winter. Some research suggests that SAD is the result of a decrease in serotonin, or an increase in melatonin, when there are fewer hours of sunlight. One treatment for SAD is therapy using bright light provided by a lamp called a light box. A daily exposure to blue light (460 nm) for 30 to 60 min seems to reduce symptoms of SAD.
In contrast, when light from a heated element passes through a prism, it separates into distinct lines of color separated by dark areas called an
atomic spectrum. Each element has its own unique atomic spectrum.
Scientists have now determined that the lines in the atomic spectra of elements are associated with changes in the energies of the electrons. In an atom, each electron has a specific energy known as its
energy level, which is assigned values called principal quantum numbers (n). Generally, electrons in the lower energy levels are closer to the nucleus, whereas electrons in the higher energy levels are farther away. The energy of an electron is quantized, which means that the energy of an electron can only have specific energy values, but cannot have values between them.
An electron can change from one energy level to a higher level only if it absorbs the energy equal to the different in energy levels.
When an electron changes to a lower energy level, it emits energy equal to the difference between the 2 levels. If the energy emitted is in the visible range, we see one of the colors of visible light. The yellow color of sodium streetlights and the red color of neon lights are examples of electrons emitting energy in the visible color range.
Each of the energy levels consists of 1 or more sublevels, in which electrons with identical energy are found. The sublevels are identified by the letters s, p, d, and f. The number of sublevels within an energy level is equal to the principal quantum number, n. For example, the first energy level (n=1) has only 1 sublevel, 1s.
The second energy level (n=2) has 2 sublevels, 2s and 2p. The third energy level (n=3) has 3 sublevels, 3s, 3p, and 3d. The fourth energy level (n=4) has four sublevels 4s, 4p, 4d, and 4f. Energy levels n=5, n=6, and n=7 also have as many sublevels as the value of n, but only s, p, d, and f, sublevels are needed to hold the electrons in atoms of the 118 known elements.
Within each energy level, the s sublevel has the lowest energy.
If there are additional sublevels, the p sublevel has the next lowest energy, then the d sublevel, and finally the f sublevel.
s < p < d < f
lowest energy < highest energy
The orbital is the 3-dimensional volume in which
electrons have the highest probability of being found.
Each type of orbital has a unique 3-dimensional shape. Electrons in an s orbital are most likely found in a region with a spherical shape.
Imagine that you take a picture of the location of an electron in an s orbital every second for an hour. When all theses pictures are overlaid, the result, called a probability density, would look like the electron could.
There are 3 p orbitals, starting with n=2.
Each p orbital has 2 lobes like a balloon tied in the middle. The 3 p orbitals are arranged in 3 perpendicular directions, along with x, y, and z axes around the nucleus.
In summary, the n=2 energy level, which has 2s and 2p sublevels consists of one s orbital and 3 p orbitals.
Energy level n=3 consists of 3 sublevels s, p, and d. The d sublevels contain 5 d orbitals.
Energy level n=4 consists of 4 sublevels s, p, d, and f.
In the f sublevel, there are 7 f orbitals. The shapes of f orbitals are complex, and we have not included them in this text.
The Pauli exclusion principle states that each orbital can hold a maximum of 2 electrons. According to a useful model for electron behavior, an electron is seen as spinning on its axis, which generates a magnetic field.
When 2 electrons are in the same orbital, they will repel each other unless their magnetic fields cancel. This happens only when the 2 electrons spin in opposite directions. We can represent the spins of the electrons in the same orbital with one arrow pointing up and the other pointing down.
An s sublevel holds one or two
electron. Because each p orbital can hold up to 2 electrons, the 3 p orbitals in a p sublevel can accommodate 6 electrons. A d sublevel with 5 d orbitals can hold a maximum of 10 electrons. With 7 f orbitals, an f sublevel can hold up to 14 electrons.
An orbital diagram shows the placement of the electrons in the orbitals in order of increasing energy.
In this energy diagram, we see that the electrons in the 1s orbital have the lowest energy level. The energy level is higher for the 2s orbital and is even higher for the 2p orbitals.
With few exceptions lower energy sublevels are filled first, and then
the "building" of electrons continues to the next lowest energy sublevel that is available until all the electrons are placed.
Chemists use a notation called the electron configuration to indicate the placement of the electrons of an atom in order of increasing
energy. The electron configuration is written with the lowest energy sublevel first, followed by the next lowest energy sublevel. The number of electrons in each sublevel is shown as a suberscript.
An electron configuration can also be written in an abbreviated configuration.
The electron configuration of the preceding noble gas is replaced by writing its element symbol inside square brackets. For example, the electron configuration for lithium 1s22s1, can be abbreviated as [He]2s1 where [He] replaces 1s2.
In period 3, electrons enter the orbitals of the 3s and 3p sublevels, but not the 3d sublevel. We notice that the elements sodium to argon, which are directly below the elements lithium to neon in Period 2, have a similar pattern of filling their s and p orbitals.
In sodium and magnesium, 1 and 2 electrons go into the 3s orbital. The electrons for aluminum, silicon, and phosphorus go into separate 3p orbitals. The remaining electrons in sulfur, chlorine, and argon are paired up (opposite spins) with the electrons already in the 3p orbitals. For the abbreviated aelectron configurations of Period 3, the symbol [Ne] replaces 1s22s22p6.
The electron configuration of the elements are also related to their position of the
periodic table. Different sections or blocks within the periodic table correspond to the s, p, d, and f sublevels.
The s block includes
hydrogen and helium as well as the elements in group 1A (1) and group 2A (2). This means that the final one or 2 electrons in the elements of the s block are located in an s orbital. The period number indicates the particular s orbital that is filling: 1s, 2s, and so on.
The p block consists of the elements
in group 3A (13) to group 8A (18). There are 6 p block elements in each period because 3 p orbitals can each hold up to 6 electrons. The period number indicates the particular p sublevel that is filling: 2p, 3p, and so on.
The d block, containing the transition elements,
first appears after calcium (atomic number 20). There are 10 elements in each period of the d block because 5 d orbitals can hold up to 10 electrons. The particular d sublevel is one less (n-1) than the period number. For example, in Period 4, the d block is the 3d sublevel. In Period 5, the d block is the 4d sublevel.
The f block, the inner transition elements, are the 2 rows at the bottom of the periodic table.
There are 14 elements in each f block because seven f orbitals can hold up to 14 electrons. Elements that have atomic numbers higher than 57 (La) have electrons in the 4f block. The particular f sublevel is 2 less (n-2) than the period number. For example, in Period 6, the f blocks is the 4f sublevel. In Period 7, the f block is the 5f sublevel.
Up to Period 4, the filling of the sublevels has progressed in order. However, if we look at the sublevel blocks in Period 4, we see that the 4s sublevels fills before the 3d sublevel.
This occurs because the electrons in the 4s sublevel have slightly lower energy than the electrons in the 3d sublevel. This order occurs again in Period 5 when the 5s sublevel fills before the 4d sublevel, in Period 6 when the 6s fills before the 5d, and in period 7 when the 7s fills before the 6d.
At the beginning of period 4, the electrons in potassium (19) and calcium (20) go into the 4s sublevel.
In scandium, the next electron added goes into the 3d block, which continues to fill until it is complete with 10 electrons at zinc (30). Once the 3d block is complete, the next 6 electrons, gallium to krypton, go into the 4p block.
Within the filling of the 3d sublevel, exceptions occur for chromium and copper.
In Cr and Cu, the 3d sublevel is close to being half-filled or filled sublevel, which is particularly stable. Thus, the electron configuration for chromium has only one electron in the 4s and 5 electrons in the 3d sublevel to give the added stability of a half0filled d sublevel. This is shown in the abbrevaited orbital diagram for chromium. A similar exception occurs when copper achieves a stable, filled 3d sublevel with 10 electrons and only 1 electron in the 4s orbital.
We can use the seasonal changes in temperatures as an analogy for periodic properties:
In the winter, temperatures are cold and become warmer in the spring. By summer, the temperatures are hot but begin to cool in the fall. By winter, we expect cold temperatures again as the pattern of decreasing and increasing temperatures repeats for another year.
The chemical properties of representative elements are mostly due to the valence electrons, which are the electrons in the outermost energy level.
these valence electrons occupy the s and p sublevels with the highest principal quantum number n. The group numbers indicate the number of valence (outer) electrons for the elements in each vertical column. For example, the elements in Group 1A(1), such as lihtium, sodium, and potassium, all have one electron in an s orbital. Looking at the sublevel block, we can represent the valence electron in the alkali metals of Group 1A(1) as ns1. All the elements in group 2A(2), the alkaline earth metals, have 2 valence electrons, ns2. The halogens in group 7A (17) have 7 valence electrons ns2np5.
A lewis symbol is a convenient way to represent the valence electrons, which are shown as dots placed on the sides, top, or bottom of the symbol for the element.
1-4 valence electrons are arranged as single dots. When there are 5 to 8 electrons, 1 or more electrons are paired. Any of the following would be an acceptable Lewis symbol for magnesium, which has 2 valence electrons.
The atomic size of an atom is determined by the distance of the valence electrons from the nucleus. For each group of representative elements, the atomic size increases going from the top to the bottom because the outermost electrons in each energy level are farther from the nucleus.
For example, in group 1A (1), Li has a valence electron in energy level 2; Na has a valence electron in energy level 3; and K has a valence electron in energy level 4. This means that a K atom is larger than a Na atom and a Na atom is larger than a Li atom.
The atomic size of representative elements is affected by the attractive forces of the protons in the nucleus on the electrons in the outermost level.
For elements going across a period, the increase in the number of protons in the nucleus increases the positives charge of the nucleus. As a result, the electrons are pulled closer to the nucleus, which means that the atomic size of representative elements decreases going from left to right across a period.
The size of atoms of transition elements within the same period changes only slightly because electrons are filling d orbitals rather than the outermost energy level.
Because the increase in nuclear charge is canceled by an increase in d electrons, the attraction of the valence electrons by the nucleus remains about the same. Because there is little change in the nuclear attraction for the valence electrons, the atomic size remains relatively constant for the transition elements.
The ionization energy is the energy needed to remove one electron from an atom in the
gaseous (g) state. When an electron is removed from a neutral atom, a cation with a 1+ charge is formed.
The attraction of a nucleus for outermost electrons decreases as those electrons are farther from nucleus.
Thus the ionization energy decreases going down a group. However, going across a period from left to right, the positive charge of the nucleus increases because there is an increase in the number of protons. Thus the ionization energy increases going from left to right across the periodic table.
In summary, the ionization energy is low for the metals and high for the nonmetals. The high ionization energies of the noble gases indicate that their electron configurations are especially stable.
An element that has metallic character is an element that loses valence electrons easily. Metallic character is more prevalent in the elements on the left side of the periodic table (metals) and decreases going from left to right across a period. The elements on right side of the periodic table (nonmetals) do not easily lose electrons, which means they are less metallic.
Most of the metalloids between the metals and nonmetals tend to lose electrons, but not as easily as the metals. Thus, in Period 3, sodium, which loses electrons most easily, would be the most metallic. Going across from left to right in Period 3, metallic character decreases to argon, which has the the least metallic character.
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