Unit 4 Atoms, Bonding, and
Chemical Reactions
Ch 18, 19, and 24
CHAPTER 18
Structure of the Atom
• Protons, neutrons,
electrons
• Quarks – small
particles that make up
protons and neutrons
Models
• Dalton - sphere
• Thompson – electrons existed
• Rutherford – nucleus containing + charge
surrounded by empty space containing electrons
• Bohr – electrons travel in orbits around nucleus
with protons and neutrons
• Electron Cloud – electrons not in fixed orbits, but
in a cloud around the nucleus
• Did The Rabbit Bite Eeyore?
Using the periodic table
• Atomic number = #
protons
• Smaller # on periodic
table
• On periodic table, #
protons = # electrons
• Atomic mass = #
protons + # neutrons
• Larger # on period
table
• # neutrons = massatomic #
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What is the atomic number of zinc?
How many electrons does tungsten have?
How many neutrons does scandium have?
What is the atomic mass of carbon?
How many protons does astatine have?
Isotopes
• Same number of
protons (same
element)
• Different number of
neutrons
• Therefore, different
atomic mass
Periodic Table structure
• Periods (left to right)
= increasing number
of protons and
electrons
• Groups (up and
down) = similar
reactive properties
Energy Levels
• 1st Level = hold max
of 2 e• 2nd Level = hold max
of 8 e• 3rd Level = hold max
of 8 e-
Drawing
• Old School Way
• Shows all the
electrons on all the
energy levels
• Ex: draw flourine
• Electron Dot diagram
• Only shows the
outermost electron
(the valence electron)
• Ex: draw flourine
Trends
• Left to right, down to up:
• Increasing electro
negativity
• Increasing ionization
energy
• Decreasing atomic radius
Cheats on your periodic table
• On your periodic table, write in the
• Group numbers
• This is the number of electrons in the
outermost level
• How many outermost electrons does
Boron have?
Bonding
• Atoms want a full outermost shell
• This is when they are most stable
• Noble Gases (far right of table) already have full
outermost shells
• Other elements want to give up or gain e- to
make a full outermost shell
• If elements lose an e-, they become positively
charged
• If elements gain an e-, they become negatively
charged
Cheats on your periodic table
Oxidation numbers
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Group 1 = becomes +1 charged ex: Li+1
Group 2 = becomes +2 charged ex: Mg+2
Group 3 = becomes +3 charged ex: Al+3
Group 5 = becomes -3 charged ex: N-3
Group 6 = becomes -2 charged ex: O-2
Group 7 = becomes -1 charged ex: F-1
• Identify the oxidation numbers for each
element:
• NaCl
• CaO
• N2O
• SiO2
CHAPTER 19
Ionic vs. Covalent bonding
• Ionic
• Total transfer of e• Between metal and
nonmetal
• On both sides of your
stairstep line
• Ex: NaCl
• Covalent
• Sharing e• Between a nonmetal
and nonmetal
• Both to the right of the
stairstep line
• Ex: CO
Identify if it is ionic or covalent
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SiO2
LiF
NaCl
C12H22O11
HCl
Polar vs. Nonpolar
• Polar
• Atoms have diff
electro negativity
• Electrons not shared
equally
• Ex: HCl
• Cl is more
electronegative then
H, therefore stronger
negative charge
• Nonpolar
• Atoms have same
electro negativity
• Electrons shared
equally
• Ex: Cl2
• Same electro
negativity
CHAPTER 24
Chemical Rxns
• Reactants --> Products
• Conservation of mass
• Mass is converted into different forms but never
created or destroyed
Symbols used in chemical
equations
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s
l
g
aq
solid
liquid
gas
aqueous, dissolved in water
Coefficients and subscripts
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4H + 02  2H20
Notice how this is balanced
Use the distributive property
4 H on left
4 H on right
2 O on left
2 O on right
Chemical Equations
• Balancing equations
• Subscripts remain the
same
• Coefficient applies to
each element
• Ex: 2HO = 2H + 2O
UNL’s tricks to balance!
• 1. Start with compound with the greatest
diversity of atoms
• 2. Leave pure elements alone until end (usually
O or H)
• 3. If rule #1 doesn’t help, start with the
compound farthest left
• 4. All coefficients must be whole numbers. This
may require multiplying by the LCM to get rid of
fraction.
• 5. # atoms of each element must be balanced
on both sides of the equation
Balance these equations
• HgO  Hg + O2
• Li + H2O  H2 + LiOH
• Mg + O2  MgO
Types of Reactions
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Synthesis A + B --> AB
Ex: 2H2 + O2  2H2O
Decomposition AB --> A + B
Ex: 2H2O  2H2 + O2
Single Displacement A + BC --> AB + C
Ex: Cu + 2AgNO3  Cu(NO3)2 + 2 Ag
Double Displacement AB + CD --> AC + BD
Ba(NO3)2 + K2SO4  BaSO4 + 2KNO3
Energy Exchanges
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Exergonic rxn = releases energy (EXITs)
Ex: glow sticks (releases light)
Exothermic rxn = releases heat
Ex: burning wood
Endergonic rxn = requires energy (moves
IN)
• Endothermic rxn = requires heat
• Ex: activating a cold pack
Catalysts vs. Inhibitors
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Catalysts
Speed up rxns
Same product is formed
Catalyst remains
unchanged and separate
from product
• Enzymes lower the
activation E, making the
rxn require less E to
occur
• Ex: enzymes break down
fruit (looks brown)
• Inhibitor
• Prevents rxn from
occurring
• Same product is formed
• Inhibitor remains
unchanged and separate
from product
• Ex: lemon juice keeps
fruit from browning