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Chapter 5

Terms in this set (47)

Periodic
Repeating pattern of similar chemical properties
Mendeleev
Developed the first (limited) periodic table of elements
It was successful, but it didn't explain why elements could be arranged by increasing atomic mass, and what was the reason for chemical periodicity?
Moseley
His work led to the modern definition of the atomic number and that the atomic number is the basis for the table's organization
Periodic law
When elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals
Periodic table
Arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same group/column
Noble gases
Most significant addition to the periodic table
Proposed by William Ramsey
Lanthanides
The 14 elements with atomic numbers from 58 to 71
They have similar physical and chemical properties
Actinides
The 14 elements with atomic numbers from 90 to 103
Periodicity
Can be observed in any group of elements in the periodic table
Valence electrons
What governs the atom's chemical properties?
S block
Chemically reactive metals
Alkali metals
Group 1
Silvery appearance
Soft
Not found freely in nature because their reactive
Alkaline earth metals
Group 2
Harder, stronger, and denser than Group 1
Higher melting points
D block
Electron configurations are not identical for different elements
Transition metals
Transition metals
High luster and conductivity
Less reactive
P block
Groups 13-18 (not He)
Valence electron number = group # minus 10
Properties vary
All 6 metalloids/semi-metals
Halogens
Halogens
Group 17
Form salts
Semi-metals
Brittle solids with some properties of nonmetals and metals
F block
Technically between Groups 3 and 4
Fill 4f sublevel
Lanthanides and actinides
Atomic radius
1/2 the distance between the nuclei of identical atoms bonded together
Orbital diagrams
Each box or line represents one orbital
Half-arrows represent the electrons
The direction of the arrow represents the spin of the electron
Shielding effect
The first two electrons in the 1st energy level shield the nucleus from the effect of 8 electrons in the 2nd level
Shielding effect
Nuclear charge is equal to atomic number
Z = atomic number = nuclear charge
(Nuclear charge is determined by the number of protons)'
S.= shielding
Bigger
The size of an atom gets _________ from top to bottom in groups because energy letters are added
Atom gets bigger
As n gets bigger, distance from the nucleus increases --> and the __________________
Same row elements
Have the same number of energy levels, but they are not the same size
Smaller (same row elements)
When the attraction is stronger, the space is condensed (they are pulled together), making the atom size _______
Number of energy levels
Take precedence over the location from left to right on the table
Ion
A charged particle
Ionization
Process resulting in formation of an ion
Cation (+)
Loses an electron
Are smaller
Loses an electron --> loses an entire energy level
Anion (-)
Gains an electron
Are bigger
Electrons pull in, but their same charge makes them repel --> attract --> repel --> attract
(Creates vibration and more repulsion)
Every electron added --> creates a greater electrical repulsion
Anion
Increase in atomic radius
Metals
Can only be positive
Generally larger atoms
Low ionization energy
Low electron affinity
Low electronegativity
Nonmetals
Tend to be negative, can be positive
Smaller atoms
High ionization energy
High electron affinity
High electronegativity
Endothermic
Any time you break an attractive force, it requires energy
Add energy to atom to break attraction with nucleus
Exothermic
Any time you form an attraction, energy is released
Ionization energy
Amount of energy required to remove an electron
Ionization energy
Endothermic
Valence electrons first, then core electrons
2nd ionization energy is always greater than the 1st
Smaller the atom, the more ionization energy
Quantum leap
The ionization energy does this when all valence electrons have been removed
Anomaly (Ionization energy)
No increase from Mg to Al
1. Instead of requiring more energy, it was less energy from 2A to 3A
2. Electron removed --> more stable
3. Because of the added stability, energy is released
4. Al may require more energy, but since it releases energy with the stability gained, then the NET change is less overall
Core electrons
Electrons in the lower levels
Electron affinity
Energy change when an atom acquires an electron
Electron affinity
Exothermic
Opposite of ionization energy
Add energy for the space of an electron
Form attraction and more stability --> NET change is higher than anticipated
Anomaly (electron affinity)
Octet
4A to 5A
(4A becomes a half filled subshell = stable
5A becomes unstable because it has 4 electrons rather than 3 in its p subshell)
Electronegativity
Ability of an atom to attract electrons within a bond
Electronegativity
Higher when the atom is smaller (valence electrons are closer to the nucleus in small atoms compared to bigger atoms)
No abnormalities, except noble gases because they don't form compounds
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