Ms. Capasso’s Awesome Unit 4 Review Sheet!
The periodic table is an organized way to display all of the known elements in the universe.
 Dmitri Mendeleev – father of the Periodic Table, arranged by increasing atomic mass
 Henry Moseley – gave us our current Periodic Table, arranging by increasing atomic number
What to know?
 Make sure you know who arranged the table by increasing mass and who arranged the table by increasing atomic number and what we use today!
Ionization Energy: The energy needed to remove one valence electron
Electron Affinity: The amount of energy released when an electron is gained
Electronegativity: The attraction an atom has for electrons in a bond
Metallic Character: How much an atom displays the characteristics of a metal
Periods - Horizontal rows running across the Periodic Table
 Each period has atoms that contain the same number of Principal Energy
Levels or electron shells
 As you go from left to right across the period, you are increasing the number of protons in the nucleus. This means the valence electrons are held closer to the nucleus by a greater force
 As you progress left to right, you start with metals and go through semimetals or metalloids and get to non-metals and then the noble gases.
What to know?
 Know which period(s) contain(s) no metals! (Hint: Look up top!)
 Know which periods contain mostly radioactive elements (Hint: Look down below!)
 Know why the trends occur; do not just memorize them!
Trends Left  Right
1.
Ionization energy increases
2.
Electronegativity increases
3.
Electron affinity increase
4.
Atomic Radii decreases

5.
Metallic character decreases
Groups - The vertical groups are called families.
 The most important thing to know about the families are that they all have the same number of valence electrons
 As you go down a family, you are increasing the number of complete electrons shells underneath the valence shell. These shells act as a shield and shield the valence electrons from the positive charge of the nucleus
Trends Top  Bottom
1.
Ionization energy decreases
2.
Electronegativity decreases
3.
Electron affinity decrease
4.
Atomic Radii increases
5.
Metallic character increases
Ionic Radii
 Cations lose electrons, get smaller
 Anions gain electrons, get larger

What you need to know?
 Which elements tend to gain or lose electrons
 How the size of their ionic radii will compare to the neutral atom
Properties of Metals, Metalloids and Non-Metals
Metals
 High melting points
 High boiling points
 Malleable – able to be hammered into various shapes
 Ductile – able to be stretched into wire
 Good conductors of electricity in all states! (Note: ionic compounds are electrolytes and are only capable of conducting electricity in the liquid state or dissolved in water! They are also brittle which helps differentiate)
 Solid at room temperature
 Have luster – shine
 Low ionization energies and low electronegativities
 Lose electrons, form smaller positive ions
Non-Metals
 Can be gases, molecular solids or network solids at room temperature
 Bromine is liquid at room temp.
 Not ductile
 If solid at room temperature, tends to be brittle (exception: network solids are very hard)
 High ionization energy and electronegativity
 Poor conductors of heat and electricity
 Tend to gain electrons forming anions that are larger than their neutral atoms.

Allotropes
 Some nonmetals can exist in two or more forms in the same phase – these are called allotropes
 Examples: O
2
and O
3
– “air” and ozone
 Diamond, coal and graphite – three forms of carbon
 STRUCTURE DETERMINES FUNCTION! If the allotropes all have different structures, they must have different functions and properties.
What to know?
 Make sure you can identify the metals, non-metals and metalloids on the
Periodic Table
 If given certain properties, make sure you can determine if that thing is a metal, non-metal or metalloid
 Be able to predict ion formations for the various groups. If I told you there was a new element found called Capasso-ium (X) and it bonded with oxygen to form this compound X
2
O, what ground would it be in? Hint: What charge would the ion have? +1….
The Families – What To Know
Alkali
 Extremely reactive!
 Do not exist in their elemental form in nature.
 Form +1 cations
 React explosively with water
 Most metallic character
Alkaline Earth Metals
 Very reactive
 Slightly less reactive than alkali metals
 Form +2 cations
 Not found in nature in elemental form
Metalloids
 B, Si, Ge, As, Sb and Te
 Display both metal and non-metal characteristics
 Semi-conductors! “Silicon Valley”
Group 14
 Carbon – important for living beings
 Silicon – an important metalloid
 Contains non-metals, metalloids and metals
 Tetravalent – can for four bonds

Group 15
 Contains non-metals, metalloids and metals
 Non-metals (N and P) form -3 anions
 N important for agriculture

Group 16
 Contains non-metals, metalloids and metals
 Oxygen is an important non-metal for living beings
Group 17
 Contains non-metals, metalloids and metals
 Exist in all three phases (F and Cl are gases, Br is a liquid and I is a solid)
 Contains only non-metals
 Held together by London Dispersion Forces
 Form -1 Anions
Noble Gases
 Stable Octet (Neon is stable with 2 electrons)
 Non-reactive
 All gases at room temperature
 Held together by London Dispersion Forces
Transition Metals
 Partially filled d-sublevel
 Hard solids with high melting point (except for mercury)
 Form ions that are colored!
Lanthanides/Actinides
 F-block
 Rare earth
What to know?
 General properties of each family
 State of elements at room temp. and the diatomic elements
 Where each family is located

Balancing Chemical Equations
Reactants: What you start with, on the left
Product: What you end up with, on the left
 The number of atoms for each element in the reactants must balance with the number of atoms in the products
 The Law of Conservation of Mass states that the mass of the reactants must equal the mass of the products.
 When balancing equations, you can only change the coefficients!
Top Five Chemical Equations
1.
Single Displacement a.
In an ionic compound, either a more reactive metal replaces the cation or a more reactive halogen replaces the anion b.
Example: NaCl + Li  LiCl + Na c.
Example 2: NaBr + Cl
2
 NaCl + Br
2
2.
Double Displacement a.
In an ionic compound, the cations of two compounds “switch partners” creating two new ionic compounds b.
Don’t forget how to recognize your polyatomic ions!!!!!!!!!
c.
Example: 2 HNO
3
(aq) + Na
2
SO
3
(aq)  2 NaNO
3
(aq) + H 2 7SO
3
(aq)
3.
Combustion: a.
Always a hydrocarbon (carbon bound to lots of hydrogen) plus O
2 producing CO
2
plus H
2
O
4.
Synthesis: a.
Small molecules or elements combining b.
Usually exothermic c.
2 Mg(s) + O
2
 2 MgO (s)

5.
Combustion: a.
A large molecule being split up into smaller molecules b.
Endothermic c.
2 NaCl (s)  2 Na(s) + Cl
2
(g)
Balancing Equations!
1.
Draw a chart with three columns – elements, # atoms reactants, # atoms product
2.
Put the symbols for the elements present in column 1
3.
Put the # of ATOMS of each element in the reactant in column 1
4.
Put the # of ATOMS of each element in the product in column 2
5.
Notice which elements are unbalanced; change the coefficient of one of the compounds or atoms above to try to balance it
6.
Change the atoms in your columns below
7.
Notice if it is still unbalanced; repeat steps 5 and 6
Helpful Hints
 Put a star next to any elements that occur in more than one compound so that you remember
 Balance the elements that occur in only one compound in p and r first; then balance the elements that occur in multiple compounds
 Balance metals first, then polyatomics, then non-metals
 Balance elements not in a compound last; these will usually be easier to balance
 If a polyatomic ion does not get broken down, try to balance it as a whole unit