Download It gets harder to take them from the atoms towards the right

Precursor Material Not in Text or Notes
1. An atom is stable (doesn’t burn, rust, react. Etc….) if it has a full s sublevel (2 electrons) and a full p sublevel
(6 electrons) – effectively having 8 electrons in its outer shell makes an atom stable - called the Octet(8) Rule
a. This electron configuration is true for the last up and down column of the periodic table (Ne, Ar, Kr….)
b. these elements are called the noble gases, they will not react with other elements, and they do not gain
or lose electrons
c. the other shell (last energy level of an atom) is called its valence shell, and its last electrons are called
valence electrons
2. All the other elements on the periodic table are trying to have the electron configuration of the noble gases (full
shells)
a. this can be done by gaining one or more electrons
(1) p5 would gain one, p4 would gain 2, p3 would gain 3….
b. or this can be done by losing electrons, and thereby stripping off 1 or 2 outer electrons to reveal a full
shell underneath
(2) s1 would lose an electron, s2 would lose 2 electrons….. and there are full p6 shells underneath
3. When elements react, they are exchanging or sharing electrons to try and either gain up to a full shell or strip off
outer electrons to expose a full shell underneath
a. if fluorine or chlorine gain one electron they have full shells
b. if sodium or potassium lose one electron they have full shells underneath
c. either process results in a stable full shell of electrons
4. For this reason, elements like sodium and chlorine react together all the time (table salt)
a. sodium loses its only 3rd energy level electron, by giving it to chlorine, and exposes its full 2 nd energy
level
b. chlorine accepts the electron in the reaction, and puts in its p6 spot to fill its 3 rd energy level
c. both are now stable atoms in a compound, which is why salt is so stable (doesn’t burn, rust…..)
5. In general elements to the left of the table try and lose electrons and elements to the right of the table try and
gain
a. the noble gases do neither, they are already full
b. Carbon and its up and down column are unique, they have 4 electrons in their outer shell, and so
sometimes they lose 4 and other times they gain 4 electrons
c. Nitrogens group usually gains 3 but occasionally can lose 5
d. Boron’s group usually loses 3 but occasionally gains 5
e. Beryllium’s group loses 2 and lithium’s group loses 1
f. Oxygen’s group gains two, and fluorine’s group gains 1
6.1 Organizing the Elements
I. The Periodic Law
A. The physical and chemical properties of the elements are periodic functions of their atomic numbers
1. In Other Words - “Elements that are in the same up and down columns, called groups or families, have
the same kinds of properties, they behave the same way ( all p 6 behave the same etc…)
B. Elements on the table are arranged in order of increasing atomic number (number of protons)
II. Metals, Nonmetals, and Metalloids – these are 3 classifications used to group different elements together by their
properties
A. Metals – found mostly to the left of the periodic table
1. Good conductors of heat and electricity
2. Lustrous (shiny)
3. Solids (except mercury)
4. Ductile (can be drawn into wire)
5. Malleable (can be hammered into thin sheets)
B. Nonmetals – found mostly to the right of the table
1. Poor conductors of heat and electricity
2. Most are gaseous
3. Solids tend to be brittle
C. Metalloids – found around the jagged red line that crosses through the p block
1. Some properties of metals, some of nonmetals
6.2 Classifying the Elements
I. Periods and the Blocks of the Periodic Table
A. Periods
1. Horizontal rows on the periodic table – side to rows
2. Period number corresponds to the highest principal quantum number of the elements in the period
Each side to side row (period) is a different Energy level on the atom
B. Sublevel Blocks
1. Periodic table can be broken into blocks corresponding to s, p, d, f sublevels
II. Blocks and Groups – blocks are the spdf sections we discussed, groups are the 18 up and down columns across the table
A. s-Block, Groups 1 and 2
1. Group 1 - The alkali metals – Just a name for the first column because they share properties
a. One s electron in outer shell
b. Soft, silvery metals of low density and low melting points
c. Highly reactive, never found pure in nature
2. Group 2 - The alkaline earth metals – a name for the second column
a. Two s electrons in outer shell
b. Denser, harder, stronger, less reactive than Group 1
c. Too reactive to be found pure in nature
B. d-Block, Groups 3 – 12
1. Metals with typical metallic properties
2. Referred to as "transition" metals
3. Group number = sum of outermost s and d electrons
C. p-Block elements, Groups 13 - 18
1. Properties vary greatly
a. Metals – some to the left of the red line
(1) softer and less dense than d-block metals
(2) harder and more dense than s-block metals
b. Metalloids – found around the red line
(1) Brittle solids with some metallic and some nonmetallic properties
(2) Semiconductors
c. Nonmetals – found to the right of the red line
(1) Halogens (Group 17) are most reactive of the nonmetals
D. f-Block, Lanthanides and Actinides
1. Lanthanides are shiny metals similar in reactivity to the Group 2 metals
2. Actinides
a. All are radioactive
b. Plutonium (94) through Lawrencium (103) are man-made
6.3 Periodic Trends – 5 properties listed below that show up on the table
I. Atomic Radii
A. Atomic Radius
1. One half the distance between nuclei of identical atoms that are bonded together
B. Trends
1. Atomic radius tends to decrease across a period due to increasing positive nuclear charge
Atoms on the left side of a row are larger and get smaller as you go to the right because the ones
on the right have more and more protons to positively attract the electrons in
2. Atomic radii tend to increase down a group due to increasing number energy levels (outer electrons
are farther from the nucleus)
There are more energy levels (rings around the nucleus) as you go down the columns
II. Trends in Ionization Energy
A. Ion
1. An atom or a group of atoms that has a positive or negative charge (by losing an electron)
B. Ionization
1. Any process that results in the formation of an ion (causes and electron to be gained or lost)
C. Ionization Energy
1. The energy required to remove one electron from a neutral atom of an element, measured in
kilojoules/mole (kJ/mol) – how much energy it takes to strip an electron off of an atom
A + energy  A + e+
D. Trends
1. Ionization energy of main-group elements tends to increase across each period
a. Atoms are getting smaller, electrons are closer to the nucleus
It gets harder to take them from the atoms towards the right
2. Ionization energy of main-group elements tends to decrease as atomic number increases in a group
a. Atoms are getting larger, electrons are farther from the nucleus
b. Outer electrons become increasingly more shielded from the nucleus by inner electrons
It gets easier to take them from the atoms on the bottom because the outer electrons are far
away
3. Metals have a characteristic low ionization energy (elements on the left are easy to steal from)
4. Nonmetals have a high ionization energy (elements on the right are harder to steal from)
5. Noble gases have a very high ionization energy (almost impossible to take an electron from last column)
E. Removing Additional Electrons
Na + 496 kJ/mol � Na+ + eNa+ + 4562 kJ/mol � Na++ + eNa++ + 6912 kJ/mol � Na+++ + e1. Ionization energy increases for each successive electron – gets harder for each next electron you take
2. Each electron removed experiences a stronger effective nuclear charge
3. The greatest increase in ionization energy comes when trying to remove an electron from a stable,
noble gas configuration
III. Trends in Ionic Size – what happens to the size of the atom if it gains or loses and electron
A. Cations – loses electrons
1. Positive ions
2. Smaller than the corresponding atom (lost electrons so the atom gets smaller)
a. Protons outnumber electrons
b. Less shielding of electrons
B. Anions – gains electrons
1. Negative ions
2. Larger than the corresponding atoms (gained electrons so the atom gets bigger)
a. Electrons outnumber protons
b. Greater electron-electron repulsion
C. Trends
1. Ion size tends to increase downward within a group
IV. Trends in Electronegativity – how much a given atom is trying to gain an electron
A. Electronegativity
1. A measure of the ability of an atom in a chemical compound to attract electrons
Once the atom is in a compound and bonded to another atom, if it is electronegative it tries to pull the
electron in the bond more to it and away from the other atom- kind of like it is greedy
Chlorine is very electronegative (greedy) when it reacts, trying to gain one more electron
2. Elements that do not form compounds are not assigned electronegativities (like the noble gases)
B. Trends
1. Nonmetals have characteristically high electronegativity
a. Highest in the upper right corner – like chlorine and fluorine and oxygen (greed to gain 2)
2. Metals have characteristically low electronegativity
a. Lowest in the lower left corner of the table – like Ra and Fr give up their electrons easily
3. Electronegativity tends to increase across a period – greedier to the right
4. Electronegativity tends to decrease down a group of main-group elements –less greedy toward bottom