Matter Chemical Trends and Chemical Bonding
Elements and the Periodic Table
-
The Nature of Atoms
o The Bohr-Rutherford model of an atom suggests that electrons exist in circular orbits
and are held by electrostatic forces with the protons in the nucleus
o Electrons have different amounts of total energy in each orbit
 For electrons to jump between orbits they absorb or emit energy in the form of
photons
 Photons are a packet of quantized energy
o Therefore, if an atom absorbs or removes energy equivalent to the level the electrons
are in it moves from a ground state to an excited state, which then emits colour when it
returns to ground state
 Called emission spectrum or line spectrum
 2 types:
o 1. Bright line spectrum
 Transition from high energy state to low causes energy
to be released as a photon of light
o 2. Dark line spectrum
 Energy is absorbed and moves from low to high
o However Bohr's model is only representative of a single electron
 So, Louis de Broglie and Schrödinger investigated the wave properties of the
energies produced by moving electrons
 This led to sublevels of orbits that only hold pairs of electrons with
opposite spins
 Therefore, each sublevel of orbits has quantum numbers which are
equivalent to the energy of each individual electron in its own
orientation of space
o A quantum number is a descriptive value that helps identify
energy level (period) and shape of orbit
o
o
A Lewis structure can help us understand the shape of a compound and how electrons
are involved in the formation of chemical compounds
Dots are placed around the chemical symbol to represent the electrons in the valence
shell
 Start by placing a single dot on top of the symbol
 Work your way around in a clockwise manor
o Only pair electrons once you get back to where you started
 Ex.
o
o
o
N

Br
1 pair of shared electrons forms a single bond, 2 pairs of shared electrons form a
double bond, 3 pairs of shared electrons form a triple bond
o Ex. N2

o

They are drawn far apart to demonstrate the repulsion that occurs between the
negative charges of the electrons
Electron pairs are less likely to be involved in chemical bonding and unpaired electrons
are trying to fill the gap

o
Ex. Draw the lewis dot diagram for the following elements:
 Mg
The shared electrons form a bonding pair and the electrons not involved
in bonding are referred to as lone pairs
We use a structural formula to connect the symbols by lines to show single, double, or
triple bonds
Rules for Lewis structure compounds:
 Position the least electronegative atom in the centre, place the other atoms
around the central atom
 Always place H or F at the end positions
o Ex. CH4
o

Ex. Draw the lewis dot diagram and structural diagram for the following
compounds; be sure to look on the P.T. to find the lowest
electronegativity vale:
o PCl3
o
o
o
o
o
Ex. CH3OH
C2H4O2
Determine total number of valence electrons (V) for all atoms
 For charged atoms be sure to add or subtract to total
o Ex. NH3, N-5 and H-1; equals 8
o EX. PO4-3, P-5, O-6, and charge +3; equals 32
Determine the total number of electrons (T) needed for each atom to have a full
electron configuration
 i.e. octet rule
Calculate shared electrons (S) in bonding; S = T – V
 Then divide by 2 to get the number of bonds
 Double bonds count as 2, and triple as 3
Calculate number of non-bonding electrons, NB = V – S
 Ex. Determine the number of bonds present in the following compounds and
draw the compound:
 CO2

C2H4

o
C4H8O3
Exceptions to Lewis Structure
 Co-ordinate covalent bonds
 1 atom contributes both electrons to the shared pair of electrons
o A charge can exist on the molecule or balanced within it
 i.e. One atom gives up an electron while another gives;
leaves the compound with 1 positively charged atom
and 1 negatively charged atom which balance each
other out
+
o ex. NH4


ex. Ozone, O3
Draw the Lewis structure for the following:
o SOCl2
o
POCl3
o
SO2Br2
o
N2O
o
H3O+

Expanded octet/valence
 Elements in third period and higher can form hybrid orbitals
o ex. PCl3 and PCl5
 therefore can hold 8 or 10
o ex. SF2 and SF6
 therefore can hold 8 or 12
 Draw the Lewis structure for the following:
o PO4-3
o
SO4-2
o
HClO3

Incomplete octet
 B and Be can form covalent bonds with halogens
o ex. BCl3 – stable with 6 electrons
o ex. BeF2 – stable with 4 electrons
 Draw the Lewis structure for the following:
o Mg(OH)2
o
o
AlP
The nucleus
 Recall: Rutherford discovered positive particles called protons in the nucleus
 He calculated the total mass of the number of protons an atom should
have to balance the negative charge of the electrons
o This mass only accounts for half the mass of a nucleus
 The other half of the mass is due to neutrons
 The chemical nature of an element depends on its atomic number
 Standard atomic notation

 Neutron number = Mass number - Atomic number
However isotopes have a different mass number and therefore a different
number of neutrons; although it is still the same chemical
 An isotope forms when energy is inputted into the atom and that
energy is used by other sub-particles to function
o When the atom returns to ground state its energy is not lost but
converted
 Therefore slight increase to mass



o
Protons are packed very closely together in the nucleus and have a repulsive
electric force between them
Protons and neutrons attract each other with a strong nuclear force which is
approximately 40 times greater than the repulsion force
 A balance between the 2, leads to a more nucleus
If neutrons and protons are not properly balanced the nucleus is unstable and it
will decay into a nucleus that is more stable
 Called radioisotopes
o Ex.
Average atomic mass
 Atomic mass of an element is usually given in atomic mass units, u
 1 u = 1.66 x 10-24 g
o Ex. Convert the mass on the periodic table to grams for the
following elements:
 Mg



Ag

Br
The unit is based on the mass of an atom of Carbon-12 and is defined as 1/12th
of the mass of a C-12 atom
 Because all other mass are compared to C-12 they are referred to as
relative masses
 We use C-12 because a smaller margin of error between the isotopes
that exist for carbon
o i.e. C-12 isotope exists naturally as 98.89%
Many elements have 2 or more natural occurring stable isotopes; an average of
their masses is taken and that is the value used on the P.T.
 Ex. Chlorine exists naturally as 75.78% Cl-35 [m = 34.97u], and 24.22%
Cl-37 [mass = 36.97 u]. What is the average atomic mass of Cl?

-
Ex. An unknown element has the following isotopic abundance: 20 % Z
[45.63 u], 35 % [25.6 u], 42 % [51.6 u], and 3 % [6.33 u]. Determine the
atomic mass and the element.
The Periodic Table
o Periodic law arranges elements by atomic number, and their chemical and physical
properties have similar observable characteristics
o Rows on a P.T. identify electrons in the same energy level
 Referred to as a period; 1-7
 Gives way to Aufbau principle that can have 4 sublevels that have
different capacities for number of electrons
o i.e. s-2, p-6, d-10, f-14
o
Columns on a periodic table are called groups
 All elements in the same group have the same number of valence electrons, and
similar chemical and physical properties
 Reactivity increases as you go down in a group
o ex. Na in water, K in water or He gas and Ar gas
o
The colours on the P.T. identify if the element is solid, liquid, or gas at standard
atmospheric temperature and pressure (which is 25 C at 1 atm)
 Also can have elements that have been synthetically produced
o
Groups on P.T.:
 Alkali Metals
 Group 1, except H
 Are very reactive, especially with water
o ex. Cesium and water is a very explosive reaction
 Halogens
 Group 17, are reactive non-metals
 Noble Gases
 Group 18, are completely un-reactive
 They do not undergo any naturally occurring reactions
 Lanthanoids
 Inner transistional metals in period 6
 Sometimes called the rare Earth elements
o ex. Some being used for battery in hybrid cars


-
Actinoids
 Inner transistional metals in period 7
 These elements have no stable isotopes which makes them radioactive
Metalloids
 Have properties between those of the metal and non-metals
 Many are shiny solids but they are poor conductors and are brittle
o ex. B, S, Ge, As, Sb, Te
Periodic Trend
o Atomic Radius
 Electrons move around the nucleus in what is best described as a cloud
 i.e. an atom has no clear define boundary
 Therefore, we define the size of an atom as the space in which electrons spend
90% of their time
 It is the distance from the centre of an atom to the boundary within which the
electrons spend 90% of their time
 For elements that can be crystallized we use X-ray crystallography
 Measure distance from the centre's of 2 atoms and the radius will be
half that distance
 Neutron diffraction and electron diffraction are used for diatomic gas molecules
 Measurements are given in picometer's, 10-12 m
 Size increases as you go from right to left on the P.T. and as you go down a
group
 Protons increase across a row and therefore electrons in orbit increase;
the charge becomes greater which intensifies the attractive forces
o Makes atom smaller
 With each step down in a group protons and electrons increase, but the
inner shells that are filled shield the outer electrons (valence) from the
positive charge of the nucleus
o The effective nuclear charge appears smaller
 Makes atom larger
o
Ionization energy
 When an atom loses an electron, the remaining ion is positively charged
 i.e. A (g) + energy → A+ (g) + e- (1st ionization)
 The amount of energy required to remove a valence electron is called the
ionization energy
 Atom and ion are in gaseous state to eliminate any effect of nearby atoms
 If solid or liquid state, the adjacent atoms would affect measurements



o
After 1 electron has been removed, it is still possible to remove more electrons,
leaving the ion with a large positive charge
 i.e. A+ (g) + energy → A2+(g) + e- (2nd ionization)
The removal of more electrons requires more energy than the removal of the
previous electron because of fewer negative charges repelling electrons in orbit
and not changing proton number
Generally increases up and to the right
Electron Affinity
 If 1 atom loses an electron, another atom must gain this electron
 Neutral atoms can gain electrons and become negatively charged
 Can release energy or may need energy to add electrons
o If energy is needed the resulting negative ion will be unstable
and will soon lose the electron
o When energy is released a stable negative ion’s formed
 i.e. A (g) + e- → A- (g) + energy
 Therefore, electron affinity is the energy absorbed or released during
the addition of an electron to a neutral atom
 Values on P.T. are negative because energy is released
 As values become more negative, they become more stable
o Unstable ions have positive values
 Generally increases as you go across a period (left to right) and up a group
o
Electronegativity
 It is an indicator of the relative ability of an atom to attract shared electrons
 Increases as you go up and right on the P.T.
 Valence shell becomes smaller across a period because the effective nuclear
charge increases pulling electrons closer to the nucleus
 If charge in EN between bonding atoms is large the atom with high EN will
attract shared electrons
 If charge in EN is small or zero the atoms will attract the electrons nearly equally
 Bond with charge in EN b/w:
 0-0.4 are non-polar covalent
 0.4-1.7 are polar covalent
 1.7-3.3 are ionic
o Polarity is how a charge is distributed throughout a compound
 Polar covalent do not share electrons equally
 Non-polar covalent almost share electrons equally
 As atoms become smaller, its EN becomes larger
o
Ex. Explain which bond would be better and why?
 Ex. 1. CaCl2, CrCl2, CoCl3

Ex.2. VF4, NbF5, TaF4

Ex. 3. NO2. PO3, As2O5
-
Ionic and Molecular Compounds
o Observable trends about elements mixing in nature:
 Metals usually form bonds with non-metals and the compounds they form are
solid
 Non-metals can bond with 1 another to form gases, liquids, or solids
 Noble gases are never found in a combined form
o Ionic bonds:
 Occur due to attractive electrostatic force between oppositely charged ions
 Therefore, atoms must be ionized before a bond can form
 Atoms gain or lose electrons to attain a filled valence shell (i.e. octet rule)
 Metals lose electrons and are positive (cations), non-metals gain electrons and
are negative (anions); which allows for the attractive force
 When forming a compound we always use the 1st charge on the P.T.; if the
compound is already formed we work backwards to determine the charge used
 Ionic compounds have an overall charge of zero
o i.e. electrons lost by metal must equal electrons gained by nonmetal
 We use a criss-cross method to form compounds
o Ex. 1. K
+ O
2. Mg + F

If charges are the same we reduce to 1:1 ratio, if numbers can be
reduced we divide by a common factor
o Ex. 1. Al + N
2. Ca + C

Recall transistional metals commonly have more than 1 charge for certain
elements
 Due to no specific pattern we need to either use the periodic table to
obtain the value or determine the possible charge from the compound
o Ex. 1. Cu + Cl
2. Fe2O3

Naming:
 Binary compounds
o Name the metal ion first, followed by the non-metal
 Name of metal is the same as name on P.T.
 If metal is a transistional metal we put a roman
numeral after the metal name that represents
the original charge of the metal used
 Name of the non-metal ion changes its ending to –ide


o
Ex.
o
Li + S
o
Al + O
o
Mn + Br
o
Nb + P
Ternary compounds
o Metal and polyatomic
o No observable pattern in naming
o We sometimes use the prefix di- to show the number of H
atoms attached to the polyatomic
o We sometimes use the prefix thio- which indicates that a sulfur
atom has taken the place of an O atom
o Name the metal first just as before, followed by the name given
for the polyatomic
o We use different prefixes and suffixes to identify the polyatomic
 Ex. ClO4- perchlorate, ClO3- chlorate, ClO2- chlorite, ClOhypochlorite
 Ex.
o Mg + HCO3
o
Al + PO4
o
Ni + SO3
Covalent bonds
 The attraction that occurs when the nuclei of 2 atoms are both attracted to 1 or
more pairs of shared electrons
 Consist of non-metal elements only and can be solid, liquid, or gas at room
temperature
 Only unpaired electrons are likely to participate in chemical bonds
 Sometimes there are not enough valence electrons for 2 atoms to share 1 pair
of electrons and form filled valence
 To complete an octet the unpaired electrons are rearranged
 When forming molecular compounds:
 Write the element that is more left on the P.T. table first (as it will act
like the metal in an ionic bond)
o



o
If in the same group write the element with the higher period
number first
Exception occurs when O reacts with a halogen; the halogen is written
first
Again use criss-cross method
o Ex. 1. N + F
2. O + Cl
3. O + S
Naming:
o The rules are exactly the same ionic bonds
 However, we assign prefixes to atoms that have
subscripts
 If the second atom is only 1 then we use the
prefix mono, this does not apply to the first
atom
o For oxygen the o, or a, are usually omitted from the prefix
 i.e. mono, di, tri, tetra, penta, hexa, hepta, octa, nona,
deca
 Ex. 1. P + O
2. C + Cl
3. N2O5
Acids and Bases
 Acids are compounds that ionize in water and release a H proton
 Bases are compounds that dissociate in water and release a hydroxide ion
 Get named like a regular ionic compound
o Ex. 1. NaOH
2. Ca(OH)2


The term alkali is often used to refer to a base that is soluble in water
 i.e. the metals of group 1 and 2
Acids in their pure form are molecular compounds
 However they are still named according to the rules for ionic
compounds with slight exceptions
o Acids that do not contain oxygen:
 Name the compound using ionic rules
 Add the word aqueous before the name
 Shows the acid was dissolved in water
 Classic old name has us remove the –ide on the
nonmetal and replace with –ic acid; begin with a prefix
hydro Ex. HCl hydrogen chloride, aqueous hydrogen
chloride, hydrochloric acid
 Ex. H2S hydrogen sulfide, aqueous hydrogen
sulfide, hydrosulfuric acid
o
-
Acids that contain oxygen:
 Are called oxoacids
 Composed of H, O and at least 1 other element (usually
a nonmetal)
 Rules for naming are the same
 However if a polyatomic ends in –ate we
change it to –ic acid
o If a polyatomic ends in –ite we change it
to –ous acid
 We no longer name the hydrogen
 Ex. HClO3 hydrogen chlorate, chloric acid; HNO3
nitric acid; HClO hypochlorous acid
Properties of Ionic and Molecular Compounds
o Melting point and boiling point
 Are unique for each pure compound
 M.P.:
o Is the temperature at which a compound changes from solid to
liquid at standard atmospheric pressure (101.325 kPa or 1 atm)
o In a solid the particles are strongly attracted to 1 another
 They are closely packed together which makes it hard to
pull apart, but they are still in motion and therefore
have kinetic energy
 Temperature is directly related to K.E. of particles
 As energy, in the form of heat, enters a
substance, the K.E. increases which increases
temperature
 When the K.E. is large enough for particles to
pull away and temperature stops increasing, the
compound will melt
o A low m.p. tells us the force between particles is small and high
b.p. has strong forces

B.P.:
o
o
Is the temperature at which a compound changes from a liquid
to a gas at standard atmospheric pressure
In a liquid particles have enough K.E. to pull away from
neighbouring particles, only to then be attracted to another
neighbouring particle
o
o
o
o
At the b.p. particles have enough K.E. to completely break away
from all other particles
A low b.p. has weak forces in the liquid and high b.p has strong
forces
Compounds with a high m.p. and b.p. are all ionic
 The attractive electrostatic forces between oppositely charged particles create
very strong bonds
Compounds with a intermediate m.p. and b.p. are molecular compounds
 Have 1 or more polar bonds
 Polar molecules have slightly negative and positive ends
o A dipole is a molecule that has the negative and positive ends
 Dipole-dipole forces are the attractive forces between
the ends of the charges
 We draw arrows to represent dipoles and if they don’t cancel out the molecule
is polar
 Ex.
o
Compound with low melting point and boiling point are all non-polar
 Because it is symmetrical all dipole forces cancel
 Ex.
o
Ex. Determine the polarity and electronegativity difference for the following
compounds. You may need to do Lewis calculagtions to get the structure.
 CH3Cl

CCl4
o

C2H6

C2H4O2

C4H4O5

H2PO4-1
It is still possible for the positive nuclei of atoms in 1 molecule to attract the electrons in
a neighboring molecule
 These attractions are very weak
 Intermolecular forces exist between all molecular compounds that aid in
attractive forces of these weaker structures
o Ex.
 dipole-dipole – based on attraction of opposite charges
(EN difference helps to observe)
 H-bonding
 Exist only between N, O, F
 Van der Waal forces (also called London Dispersion)
 Sum of attractive and repulsive forces between
molecules
 Attraction between dipole of molecule and
dipole of surrounding molecules

Intramolecular forces are forces that exist within compounds
 Due to the transfer of electrons or sharing of electrons
o
Solubility in water
 For a substance to dissolve in water, the water molecules must be more strongly
attracted to particles of the substance than to other water molecules
 Water is polar and its positive end will attract a negative ion or negatively
charged end of another polar molecule
 Negative end of water attracts positive ion or positively charged end of
another polar molecule
 Therefore, water is likely to dissolve ionic and polar compounds
 LIKE DISSOLVES LIKE
o Ex. Sucrose is polar and table salt is ionic; both are soluble in
water
 Water molecules are more strongly attracted to each other than to non-polar
molecules
 Therefore do not dissolve in water
o Ex. Fats and oils
 Note not all ionic compounds and polar molecular compounds are soluble in
water
o
Electrical conductivity
 Is the ability of a substance to allow an electric current to exist within it
 A substance can conduct an electric current only if charges (electrons or ions)
can move independently of 1 another
 A pure metal allows electrons to flow freely because they are not tightly bound
to the metal atoms
 When conducting electric current, electrons are moving with ease from
1 metal atom to the next
 A pure ionic compound can only conduct an electric current under conditions in
which entire ions can move independently of 1 another
 In their solid form ionic compounds cannot conduct an electric current
o Can conduct in liquid state or aqueous state


In molecular compounds electrons never leave 1 atom completely
 Therefore, there are no positive and negative charges that are
independent of 1 another
o Therefore, cannot conduct electric current
Recall acids are molecular compounds but have ionic properties when dissolve
in water; therefore can conduct electric current