UNIT 1 – STRUCTURES AND PROPERTIES OF MATTER
Developing a Nuclear Model of the Atom o The model of an atom has drastically changed from a lawn-bowling ball, to plum pudding, to a planetary system, to a quasi-particulate entity surrounded by even smaller fundamental entities that are neither particles nor waves

Thomson first discovered the presence of electrons through the use of a cathode ray tube which sent an electric current from the cathode to the anode and that energy stimulated the gas particles in the tube and gave off a colour

Rutherford tested this theory by shooting alpha particles (which are positively charged) at a piece of gold foil with a zinc sulfide coated screen near the gold o The alpha particles bounced off at odd angles leaving him to believe electrons must be travelling around a core of positive charges that he called the nucleus

Deflections caused by intense electric field

However if his model was correct

An electron in motion around a central body MUST continuously emit radiation o Therefore, we should see a continuous spectrum of light energy

Also because electrons should lose energy the radius of the orbit should decrease spiralling into the nucleus obliterating the atom entirely

OR Rutherford did not have the whole picture yet o Light is a form of energy called electromagnetic radiation

Energy transmission in which electric and magnetic fields are propagated as waves through empty space (vacuum) or through a medium

Ex. radio waves, microwaves, infrared, visible, UV, x-rays, gamma rays o As wavelength decreases, frequency increase o Electromagnetic waves come in contact with matter as packets or quanta of energy called photons

The energy of a photon depends on the wavelength or frequency

Ex. long wavelength, have low frequency and lower energy photons o λ = c/f
 where c = 3.0 x 10 8 m/s

Planck discovered a relationship between the energy of a photon and its frequency o E = h/f

Where h = 6.626 x 10 -34 Js (Planck’s constant)

o Finally Bohr came along and using Planck’s and Einstein’s previously designed theories determined the following:

Electrons exist in circular orbits and are held by electrostatic forces with the protons in the nucleus
 Electrons can exist only in “allowed” orbits

Have different amounts of total energy in each orbit o Means that energy of electrons in atoms is quantized

While electrons remain in 1 orbit it does not radiate energy

Electrons can jump between orbits (energy levels) by absorbing or emitting photons

Must have the same amount of energy as the difference in the energy levels of the electrons 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

If this colour produced passes through a prism, dispersion occurs emitting a specific set of colours related to that element

Called emission spectrum or line spectrum

Bright line spectrum - transition from high energy state to a lower causes energy to be released as a photon of light

Dark line spectrum – energy is absorbed and moves from low to high energy state o When electrons are attracted to the nucleus and confined to the orbit, the electron energy becomes negative

En = -R
H
/n 2


Where R
H
= 2.179 x 10 -18 J o As energy approaches zero, the electron can be in an infinite number of orbital’s
 Ionization energy also increases
 Bohr’s model is only representative of a single electron in orbit

His theory also had difficulty explaining the effect of magnetic fields
Subatomic particles: o Protons – positively charged particles with a relative mass of 1 o Located in the nucleus o Electrons – negatively charged particles with a relative mass of approximately 1/1836 o Found travelling in regions of space around the nucleus o Neutrons – neutral particles with a relative mass of 1 o Located in the nucleus o Quarks – make up the protons and neutrons o Gluons – stick the quarks together o Photons – light is made of these particles o Leptons – make up electrons and particles like electrons o Neutrinos – are neutral particles produced in nuclear reactions

Decays into a proton and electron o Bosons – carriers of the weak nuclear force o Hardons – have a strong interaction that assists in forming a lot of the particles

In the end there are hundreds of subatomic particles and new ones always being discovered
The Quantum Mechanical Model of the Atom o Louis de Broglie suggested light interacts with matter as individual photons

Therefore, reasoned that light can behave as a wave and particles based off observable properties
 λ = h/mv o Ek = hf (kinetic energy – energy of motion)

He concluded that the radius of the orbit would equal an integer number of wavelengths or the waves would cancel out
 Helped back Bohr’s model o Schrodinger later produced a wave equation to explain elements with more than 1 electron



The wave function contains 3 variables that are called quantum numbers
 n, l, m l
When you sub specific combinations of integer values for the variables you get different solutions to the wave equation

Each solution describes a region of space around the nucleus of an atom o Each quantum number represents a specific characteristic of the electron occupying that space

The quantum mechanical model describes electrons treated as waves

Heisenberg developed an uncertainty principle that could help determine the probability of finding an electron within a region of space described by the wave equation
 ∆x∆mv ≤ h/4π o Where x is the position of the particle, mv is the momentum, and h/4π = 5.2728 x 10 -35 Js

Therefore an atomic orbital appears more as a fuzzy cloud in which its density at any point is proportional to the probability of finding the electron at that point
 i.e. like throwing darts at a dart board o Quantum numbers describe electrons in atoms; 3 types

First quantum number, n

Describes distribution of electron in atom
 Determines specific energy level (shell) of atomic orbital’s and its relative size
 All orbital’s that have the same value of n are said to be in the same sell
 n can range from 1 to infinity


 high n value refers to high energy level/shell
Second quantum number, l

Refers to the orbital shape

Refers to energy sublevels (or subshells) within each principal energy level

Values are dependent on n

Values are positive integers that range from 0 to (n-1)
 Therefore, if n =2; l = n-1 = 1 o If l=0 it is s orbital o If l = 1 it is p orbital o If l = 2 it is d orbital o If l = 3 it is f orbital
 Depict the shape and orientation of density region of space in which electrons can be found

To identify energy sublevel combine n and l o Ex. n = 3, l = 2 --- 3d

Magnetic Quantum number, m l

Indicates the orientation of the orbital in the space around the nucleus

Depends on the value of l o m l
= 2l + 1

Ranges from –l to +l o Ex. if l = 0, m l
= 0; m l
= 2l + 1 = 1
 If l = 1, m l
= -1, 0, +1; m l
= 2l + 1 = 3

Spin Quantum number, m s

Electrons spin on their axis and the spinning charge generates a magnetic field

Specifies direction in which axis of electron is oriented and only has 2 possible values: + ½ or – ½

A pair of electrons with opposing spins has no net magnetic field
 + ½ is represented by ↑ and spins clockwise, - ½ is represented by ↓ and spins counter clockwise

o The Pauli exclusion principle states that only 2 electrons of opposite spin could occupy an orbital

Therefore, an orbital can only have a maximum of 2 electrons which must have opposite spins

Can have 1 electron with either spin or no electrons in the orbital

Each electron in an atom has its own unique set of 4 quantum numbers

Ex. if n is considered a city, in the city there is a finite number of streets, each street is represented by l, each street has a finite number of buildings each with its own number which can be represented by m l
, finally the name of the person at the address is represented by m s o Ex. lithium --- electrons – 3
 Electron 1 – n=1, l= 0, m l
= 0, m s
= + ½

Electron 2 – n =1, l= 0, m l
= 0, m s
= - ½

Electron 3 – n = 2, l = 1, m l
= -1, 0, +1, m s
= + ½ o Ex. What are n, l, and ml for a 5d orbital?
 n= 5, l= 2, m l
= -2, -1, 0, +1, +2
 For a 3f orbital?
 n= 3, l =3, m l
= -3, -2, -1, 0, +1, +2, +3
Electron Configuration and the Periodic Table o Electrons in different orbital’s within the same energy level will have different energies

If only 1 electron then only attractive forces exist; if more than electron then both attractive and repulsive forces exist o Regardless of energy absorbed and electron excitation there will be no difference in the energy of 1 electron whether it is in s, p, d, f for any specific n
 With more electrons s, p, d, f orbital’s have different energies for any specific n

Can be seen in an energy level diagram o s orbital’s are closer to the nucleus, and therefore require less energy
 Keep in mind all orbital’s within the same sublevel have the same energy
 Ex. 2px, 2py, 2pz o s orbital’s have 1 plane o p orbital’s can be found in x, y, z planes o d orbital’s can be found in x 2 -y 2 , xy, xz, yz, and z 2

An atoms electron configuration shows the number and arrangement of electrons in its orbital’s
 Ex.

o Orbital diagrams:

A box with no electrons is a unoccupied orbital

A box that is half filled has 1 electron

A box that is filled has 2 electrons in opposite directions, i.e. spin

Ex. 1s 2

Can only have a max of 2 electrons o Aufbau principle states that each electron occupies the lowest energy orbital available

Fill orbitals of the same energy level before going to the next orbital

When electrons are added in the same energy sublevel each orbital gets 1 electron before pairing occurs o Electrons in different orbitals of same energy have the same spin
 Hund’s rule:

Single electrons with the same spin must occupy each equal energy orbital before additional electrons with opposite spins can occupy the same orbitals
 s block elements group 1 and 2 (max 2 electrons)
 p block elements group 13-18 (max 6 electrons)
 d block elements group 3-12 (max 10 electrons)
 f block elements are for inner transitional metals (max 14 electrons) o Ex. Mg – 12 electrons

1s 2 2s 2 2p 6 3s 2

Superscripts add to 12 o Ex. Br – 35 electrons
 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5


Or we can use a previous period noble gas to give a short form o Ex. argon – 1s 2 2s 2 2p 6 3s 2 3p 6
 Due to its stability and no charge

[Ar]4s 2 3d 10 4p 5

For ions:

It is easier to remove electrons from lower energy orbitals or orbitals that are not complete

When adding electrons we fill orbitals or move to the next highest energy level o Ex. N – 7 electrons, charge of -3

1s 2 2s 2 2p 3 ---- becomes 1s 2 2s 2 2p 6 o Ex. Pb – 82 electrons, charge of +4

[Xe]6s 2 4f 14 5d 10 6p 2 ---- becomes [Xe]4f 14 5d 10 o Periodic table trends

Atomic Radii

Determine size of an atom by measuring distance between the nuclei of adjacent atoms o For metals: between atoms in a crystal; for molecules: between atoms bonded together

Increases as you go left and down the periodic table o However reduced force of attraction for larger radii due to distance of electrons from nucleus
 The reason it increases left is because as you go right in period ore protons are added, therefore increasing the force of attraction causing electrons to be closer

Pattern not as observable in transitional metals due to the d-orbital electrons (shape)

Ionization energy

Energy required to remove an electron from a ground state atom in a gaseous state o To remove an electron energy is needed to overcome the force of attraction

The first ionization energy is the least amount of energy required to remove an electron from the valence shell o Is closely linked to its chemical reactivity

Low IE usually from cations (group 1); high IE usually from anions (group 17)


Opposite property to atomic radius o Increase as you go right and up the periodic table o Some variations in IE with B, Al, O, S because by removing 1

Electron affinity electron it can make for a stable filled or half filled orbital

Change in energy that accompanies the addition of an electron to an atom in the gaseous state o Addition of 1 mol of electrons to 1 mol of gaseous atoms/ions

First electron affinity results in an anion o Energy is released when the first electron is added because it is attracted to the atoms nuclear charge

Second electron affinity is always positive because energy must be absorbed

Large negative numbers mean high EA; low negative numbers and positive numbers mean low EA
Models of Chemical Bonding o Electronegativity

Ability of atoms to attract shared electrons in a chemical bond

Increases as you go up and right on periodic table

The valence shell becomes smaller across a period because the effective nuclear charge increases pulling electrons closer to the nucleus
 If ∆EN between bonding atoms is large the atom with high EN will attract shared electrons
 If ∆EN is small or zero the atoms will attract the electrons nearly equal o Bonds with ∆EN between:

1.7 – 3.3 are mostly ionic

0.4 – 1.7 are polar covalent

0 - 0.4 are mostly covalent (also called non-polar covalent)

Not an accurate scale as there may be variations

o Ex. hydrogen halides (HF, HI, HCl, HBr) range from 0.4-1.9 and behave like covalent bonds but can ionize in water to form strong acids

Ex. MgO ---- ∆EN = 3.4 – 1.3 = 2.1; therefore would be ionic o Types of bonding:

Metallic Bonding

Have low EN and valence shells that are less than half filled

Have trouble attracting and holding electrons of other atoms well enough to form filled valence shells

Metal atoms attract their own electrons so weakly that the valence electrons of solid or liquid metals can move somewhat freely to other atoms
 Electrons are said to be delocalized because they don’t remain in 1 location or in association with one specific atom o Described by an electron sea model
 The model is an array of cations in a “sea” of freely moving electrons with the positively charged ions attracted to many of the electrons in the “sea” simultaneously

Atoms of metals form uniform crystalline patterns o Atoms of the crystal form precise, regularly repeating patterns but atoms at the boundaries of the crystal are arranged randomly

Some properties of metals include: m.p., b.p., electric and thermal conductivity, malleability, ductility, and hardness

Alloys are the most commonly used metals o Mixture of 2 or more different types of metal atoms
 Substitutional alloy

If atoms are the same size it can take the place of an atom in a pure metal

Interstitial alloy

If atoms are smaller than the pure metal they may fit between the larger pieces

Ionic Bonding

Only found between interactions of metals and non-metals

Although some sharing occurs, 1 atom essentially loses 1 or more valence electrons and the other atom gains

Ionic crystals form due to the attraction of the positive and negative atoms trying to be as close as possible o i.e. crystal lattice

 for ionic compounds to be soluble in water the attractive forces between ions and water molecules must be stronger than the attractive forces on the ions themselves o ex. salt in water
 the higher the m.p. the less likely it is for an ionic compound to dissolve in a solvent

Covalent Bonding

Sharing of electrons between atoms of non-metal elements so that both valence shells are complete

2 types: polar and non-polar o Polarity is how a charge is distributed throughout a compound o Polar covalent bonds do not share electrons equally
 Electrons tend to be near the atom with higher EN o Non-polar covalent bonds almost share electrons equally

Negatively charged atom is labelled, ᵟ
-

Positively charged atom is labelled, ᵟ
+

Show bond polarity by placing an arrow from positive to negative end

The nuclei of both atoms exert attractive forces on the shared electrons o Distance and shape are determined by repulsive forces between nuclei and between electrons

A network solid is a substance in which all atoms are covalently bonded together in a continuous 2 or 3 dimensional array; no natural beginning or end exists o Allotrope – 1 of 2 or more compounds consisting of the same element but having different physical properties
Shapes, Intermolecular Forces and Properties of Molecules o Each atom, bonding pair of electrons, and lone pair of electrons takes up space and has its own position in space relative to other atoms and electrons

Different forces of attraction and repulsion determine these positions

o Ways to predict molecular shape:

Lewis Structures

Allows you to see how many electrons are involved in each bond o Also how many lone pair electrons are present
 Gives a 2-D image

Steps for drawing o Position the least EN in the centre, place the other atoms around the central atom and draw a single bond between each

Always place H or F at the end position o Determine total number of valence electrons (V) for all atoms
 For polyatomic ions be sure to add or subtract charge o Determine the total number of electrons (T) needed for each atom to have a full electron configuration
 i.e. octet rule o Calculate shared electrons (S) in bonding. S = T-V

We divide by 2 to get number of bonds

Double bonds count as 2, triple as 3 o Calculate the number of non-bonding electrons (NB). NB= V-S
 Add these electrons to the atom to fill the electron configuration

Exceptions in Lewis structures o Co-ordinate covalent bonds

One atom contributes both electrons to the shared pair of electrons

Ex. NH
4
+ o Expanded octet/valence
 Atoms of the third period and higher can form hybrid orbitals

Ex. SF
6

o Incomplete Octet
 Boron and beryllium can form covalent bonds with halogens

Ex. BCl
3

Ex. BeF
2

Examples:
1. NF
3 o Resonance structures

Show same relative position of atoms but different positions of electron pair

Ex. O
3

Ex. CO
3
-2

2. PO
3. CCl
4
-3
2
O

VSEPR Structures (valence shell electron pair repulsion theory)

A molecules shape affects its physical and chemical properties

Bond length and angle determine the shape o Angle is formed by 2 atoms that surround the central atom

VSEPR theory states that electron groups around an atom are positioned as far as possible from the others to minimize repulsion o An electron group is a single, double, or triple bond, or lone pair

Shapes

A is central atom, X is other atoms, E is lone pair o Linear, AX
2

Bond angle of 180 o Trigonal Planar, AX
3
 Bond angle of 120

V-shaped or bent, AX
2
E

Bond angle less than 120

o Tetrahedral, AX
4
 Bond angle of 109.5

Trigonal pyramidal, AX
3
E

Bond angle less than 109.5
 V-shaped or bent, AX
2
E
2

Bond angle less than 109.5 o Trigonal Bypyramidal, AX
5

Bond angle of 90 up and down, 120 on the sides
 Seesaw, AX
4
E

Bond angles less than 90 up and down, less than
120 on the sides

T-shaped, AX
3
E
2

Bond angle less than 90


Linear, AX
2
E
3

Bond angle of 180 o Octahedral, AX
6

Bond angle of 90
 Square Pyramidal, AX
5
E

Bond angle less than 90
 Square Planar, AX
4
E
2

Bond angle of 90 o Hybridization

Valence bond theory suggests that a covalent bond forms when the atomic orbital’s of 2 atoms overlap to share a common region in space and a pair of electrons occupies that region overlap
 Molecular orbital theory states that when atomic orbital’s overlap they combine to form new orbital’s called molecular orbital’s o Have new shapes and energy levels

Electrons are delocalized throughout the new orbital

Orbital overlap has max capacity of 2 electrons that have opposite spin o i.e. Pauli Exclusion Principle

The greater the orbital overlap the stronger and more stable the bond

Types of hybridization:

Single bond o When a new molecular orbital forms the bond is called a sigma bond, ϭ
 It is a bond that is symmetrical around the bond axis of the two nuclei

Which means you can rotate the bond around the bond axis and nothing will change

o A hybrid orbital is formed by the combination of 2 or more orbitals in the valence shell of an atom
 1s and 3p orbitals combine to form sp3 orbitals

Ex. CH
4
– 1s 2 2s 2 2p 2

Double bonds o A double bonded compound contains a sigma bond and a pi bond, π
 The π bond is constantly shifting in space around the compound and because of this movement they are easier to break

1s and 2p orbitals combine to form sp2 orbitals o Ex. C
2
H
4

Sometimes may draw pi bonds as lobes around the molecule showing the region of space the electrons in the pi bond have to move


Triple bonds o P orbitals exist in 3 planes; x, y, and z which therefore allows for the possibility of triple bonds
 1s and 1p orbital combine to form sp orbitals

Ex. C
2
H
2

General shapes and names for hybrid orbitals

Linear, sp

Trigonal planar, sp 2

Tetrahedral, sp 3

Trigonal bipyramidal, sp 3 d

Octahedral, sp 3 d 2

o Influence of molecular shape on polarity

Recall a covalent bond is polar when 2 atoms have an EN difference between
0.4-1.7

A partial negative charge on 1 atom and a partial positive charge on the other atom form a bond dipole o Ex. H-F

Polarity can be determined by adding the vectors o Identical polarities around a central atom tend to form non-polar molecules

Ex.

H
2
O

CO
2

CCl
4

CHCl
3 o So far we have examined intramolecular forces in covalent bonds

Force that holds atoms or ions together

Influence chemical properties of substances o Chemical changes involve overcoming these forces in order for bonds to break and new substances to be synthesized o Intermolecular forces

A force that exerts between molecules or between ions and molecules to influence the physical properties of substances

It is the force of attraction and repulsion between molecules


Types:

Dipole-Dipole o The partial positive and negative charges of a polar molecule form permanent dipoles which are always present in the molecule
 Neighbouring molecules align themselves so that oppositely charged regions are directed toward each other

Strength depends on polarity o Main reason for difference between m.p. and b.p.

Hydrogen bonds o Are dipole-dipole interactions

When covalently bonded to a high EN atom it draws the electrons away from the H o i.e. N, O, F
 lone pairs on EN atom enhance attractive forces

Ex. Why is water a liquid at room temperature, while ammonia is a gas?

Both are similar size, polar, and have H bonding

However forces holding ammonia together are weaker because N is less EN than 0

N-H bonds are less polar than O-H; therefore weaker H bonds

N has 1 lone pair and O has 2

A single O atom of 1 water molecule can be Hbonded to as many as 6 H atoms at different water molecules at the same time; ex. gives snowflakes their characteristic six-sided shape

Ion-Dipole o Ions and dipoles are attracted to 1 another by electrostatic forces o 2 types:

Polar molecule and a cation

Polar molecule and a anion o Strength depends on charge and size of the ion and on magnitude of dipole and size of molecule
 If the charges have the same magnitude the cations will

Induced Dipoles interact more strongly with dipoles o Dipole induced dipole forces is when the charge on a polar molecule is responsible for inducing the charge on the non-polar molecule o Ion induced dipole force is when a nearby ion can induce a dipole in a non-polar molecule


Dispersion forces o A weak attraction between all molecules , including non-polar, due to temporary dipoles
 Therefore only force in non-polar molecules o Attraction becomes larger as the mass of the molecule increases

Greater number of electrons also increases force strength

Probability of temporary dipole forming increases as electron number increases
 Shape of the molecule also effects the strength