The nucleus is made up of protons (p+) and neutrons (no) and is surrounded by rings of orbiting electrons (e-)

Subatomic particle

Relative mass Relative charge

Proton 1 +1

Electron 0 -1

Neutron 1 0

Also called isotope notation

X = element symbolA = atomic mass = # protons + # neutronsZ = atomic number = # protons

XAZ

What is the difference between these two atoms?

Introduction: BONDING Chemical reactions involve the formation of

new products Bonds between atoms or ions in the

reactants must be BROKEN (this requires energy, and is therefore an ENDOTHERMIC process)

Bonds are then FORMED between atoms or ions to make the products of the reaction (this releases energy, and is said to be an ENDOTHERMIC process)

Intramolecular vs Intermolecular Forces The bonds that are present between atoms within a

molecule are due to INTRAmolecular forces of attraction – relatively STRONG forces of attraction that hold the molecule together

Forces of attraction between molecules, INTERmolecular forces of attraction, are what keeps molecules together as a solid, or liquid;

INTERmolecular forces of attraction must be overcome if a substance is to change its physical state from solid to liquid to gas: The strength of these intermolecular forces is affected by the type of bonding within the molecule

o We will begin with a study of the types of bonding found in compounds, and then look at the characteristics of these bonds and how it affects the PHYSICAL PROPERTIES (like solubility, melting point, boiling point, conductivity and hardness) of different compounds

Ions are elements that have gained or lost electrons

Ions are commonly found dissolved in water, such as in the cytoplasm or plasma of the blood

Elements in the same family tend to form the same type of ion (e.g.: Na+, Li+, K+, Rb+)

Some important ions are Ca2+ (used for muscle contraction), Na+ and K+ (nerve and muscle function), Fe2+ and Fe3+(in hemoglobin) and H+ (required for synthesis of ATP)

Electrons orbit the nucleus of an atom at a great distance compared to the size of the particles

Analogy: If an apple represented the size of an atom’s nucleus and it was placed at the center of the earth’s core, the valence electrons would be orbiting close to the surface of the earth’s crust

The valence electrons therefore are the part of the atoms that interact in chemical reactions to form compounds

Ionic Compounds form when electrons are transferred from one atom to another to form ions with complete outer shells of electrons (ie the same stable electron configuration as the inert gases.)

Ionic BONDS form between a metal and a nonmetal eg NaCl

Metal tends to lose electrons which are transferred to the nonmetal

Metals form a cation (+) and nonmetals form an anion (-)

Formation of NaCl Formation of MgF2

Ionic Compounds

Positive and negative ions are held together in a strong lattice framework through the strong electrostatic attraction of oppositely charged ions

Ionic compounds have high melting points and are solids at room temperature

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.prenhall.com/wps/media/objects

These result in a lattice of ions rather than individual molecules, so we refer to MgF2

and NaCl as formula units, not molecules. Properties of ionic substances as a result of

the strong, rigid lattice that can dissociate when dissolved in water:◦ Crystalline solids at room temperature◦ Hard and brittle◦ High melting and boiling points◦ Conduct electricity when in liquid form◦ Most are soluble in water

Ionic Compounds are crystalline:

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Dissolving of NaCl in water: http://www.youtube.com/watch?v=EBfGcTAJ

F4o&feature=related

http://www.youtube.com/watch?v=dr4sFNzUVzI&feature=related

NaCl solutions conduct electricity http://www.youtube.com/watch?v=aELPrWzi

xeU&feature=related

The chemical formula of simple ionic compounds can easily be determined if one knows how many electrons an atom needs to lose or gain to achieve a stable inert gas electron configuration – this is related to its group number = # of valence electrons already present

In an ionic compound, the sum of positive and negative charges = zero to create a neutral compound e.g. MgS, Al2O3 CaBr2

Group # indicates charge for families 1,2,3 and 5,6,7

1 2 3 4 5 6 7

+1

Li

+2

Be

+3 -3

N

-2

O

-1

F

Na Mg Al P S Cl

Formulas and Names of Ionic Compounds Simple Ionic Compounds: metal + nonmetal

Eg sodium and bromine

Names and Formulas of ionic compounds continued……. Calcium and chlorine

Simple Ionic Compounds To name, state the name of the metal

element first, the nonmetal element second and change the ending of the nonmetal element to “ide”. No prefixes are used.

eg NaCl sodium chloride

Eg MgBr2 magnesium bromide

Words formula Use the crossover rule to determine

the formula, using the valence or ionic charge of each ion. The resulting compound should be neutral (net charge = zero)

Eg aluminum oxide eg lithium bromide

For ionic compounds, always reduce the ratio of metal : nonmetal to lowest terms

Eg Magnesium and oxygen

Alternatively, use an accounting method instead of the crossover: Calcium and oxygen or calcium and

chlorine

Multivalent Metal Ions

Iron (II) Fe2+ FeCl2 iron (II) chloride

Iron (III) Fe3+ FeCl3 iron (II) chloride

Copper (I) Cu1+ Cu2O copper (I) oxide

Copper (II) Cu2+ CuO copper (II) oxide

Tin (II) Sn2+ SnO tin (II) oxide

Tin(IV) Sn4+ SnO2 tin (IV) oxide

Lead (II) Pb2+ PbO lead (II) oxide

Lead (IV) Pb4+ PbO2 lead (IV) oxide

Polyatomic ions Some ions are composed of a number of

nonmetal elements bonded together covalently, with an overall net charge due to a surplus or deficit of electrons

Eg NH4+ the ammonium ion

Eg SO42- the sulphate/sulfate ion

Common Polyatomic ions

NH4+ ammonium SO4

2- sulfate

OH- hydroxide SO32- sulfite

NO3- nitrate HSO4

2- hydrogensulfate

NO2- nitrite PO4

3- phosphate

ClO3- chlorate HPO4

2- hydrogen phosphate

CO32-

HCO32-

CarbonateHydrogen carbonate

H2PO41- dihydrogen

phosphate

Memory Aid for Polyatomic Ions “Nick the Camel ate a Clam for Supper in

Phoenix”

First letter: the non metal element

# consonants = # of O’s in formula

# vowels = # of negative charges on the ion OR # of hydrogens in the related oxyacid

Example: Nick “N” is for nitrogen N

3 consonants = 3 O’s NO3

1 vowel = 1 negative charge NO3 –

or 1 H in the oxyacid: HNO3(aq)

Now you try……. Camel tells us:

Nick the Camel gives us the –ate polyatomic ions, aka the “MOTHER IONS” (by Mrs. Wheelihan)

When 1 oxygen is removed from the oxyion, the ending changes from -ate to –ite, BUT THE CHARGE STAYS THE SAME

EG NO3- nitrate and NO2

- nitrite

SO42- sulfate and SO3

2- sulfite

This idea can be used to name a number of variations from the “mother ions”

Sometimes a hydrogen ion, H+, stays attached to a polyatomic ion This creates a new polyatomic ion with the

prefix “hydrogen” or “dihydrogen” Eg hydrogen carbonate ion

Other examples:

PO43- phosphate

HPO42- hydrogen phosphate

H2PO41- dihydrogen phosphate

NaH2PO4 sodium dihydrogen phosphate

Peroxide Ion, O22-

PEROXIDES contain the peroxide ion, O22-

instead of the oxide ion, O2-

example, hydrogen peroxide H2O2

Peroxides are named by placing the prefix “per” before the word oxide in the name of the compound. When writing the formula of a peroxide, simply add an additional oxygen atom to the formula of the normal oxide, representing the peroxide ion, O2

2-

Eg lithium peroxide Li2O2

Form between two nonmetals Electrons are shared rather than

transferred Macromolecules and organic molecules

are covalent molecules held together with covalent bonds.

Examples: lipids, carbohydrates, proteins and nucleic acids.

What holds a covalent bond together?

Nuclei repel each other, but are both attracted by the pair of negative electrons being shared.

Formation of H2 Formation of NH3

Other examples, with their Lewis diagrams:

Formation of O2 – a double bond

“Lone pairs” and “Shared pairs”

Non-bonded pairs are called “LONE PAIRS”

Pairs of electrons that are shared between atoms are called: ”BONDING PAIRS”

BOTH LONE PAIRS AND BONDING PAIRS MUST BE SHOWN ON YOUR LEWIS DIAGRAM

ie ALL VALENCE ELECTRONS must be shown

Bonding pairs and Lone Pairs

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Lewis Structures of Covalent Molecules Covalent molecules can be represented in a

number of different ways. In Lewis structures, ONLY THE VALENCE ELECTRONS ARE SHOWN.

In Lewis structures, the outside electrons are shown with dots and covalent bonds are usually shown by bars, but there are several ways of drawing Lewis structures (see Figures 5 to 9, Neuss (2007) pages 64, 65)

There are a number of different ways to represent Lewis structures of covalent molecules:

These can also be drawn with lines or crosses to represent lone pairs and bonding pairs of electrons

. Covalent Bonds usually involve the sharing of an electron pair consisting of 1 electron from each molecule involved in

the bond

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Coordinate or Dative Covalent Bonds

Sometimes the “bonding pair” of electrons both come from the same element. This type of bond is called a “coordinate or dative covalent bond”.

Eg ammonia

Chewtychem.wiki.edu

Other examples of dative covalent bonds (also called coordinate bonds) are found on page 65 of Neuss (2007) and include carbon monoxide, CO(g); the hydronium ion, H3O+; and aluminum chloride dimer

Single Covalent Bond When 1 pair of electrons is shared between

the same two atoms, a SINGLE COVALENT BOND is formed

Double Covalent Bonds When 2 pairs of electrons are shared

between the same two atoms, a DOUBLE COVALENT BOND is formed as found in CO2 (g), carbon dioxide or carbon (IV) oxide

Carbon dioxide CO2 http://dbhs.wvusd.k12.ca.us

Triple Covalent Bonds When 3 pairs of electrons are shared

between atoms, a triple bond is formed Eg, nitrogen gas, N2

http://www.webchem.net/notes/chemical_bonding/covalent_bonding.htm

BOND LENGTH and BOND STRENGTH

The more pairs of electrons shared between atoms, the stronger and shorter the bond

Therefore single bonds tend to be longer and weaker than double bonds and triple bonds are shorter and stronger than corresponding double bonds, as shown in Table 1 below

Table 1: Bond length and bond strength (Neuss 2007 p. 66)

Example Length (nm) Strength (kJ mol-1)

O-O

Cl-Cl

N-N

C-C

H-H

0=0

N=N

C=C

C≡C

N≡N

If single, double, and triple C to C bonds are compared a distinct pattern emerges:

Bond Bond strengthkJ/mol-1

Bond length

nmC-C 348 0.154

C=C 612 0.134

C=C 837 0.120

Describe the trend in bond strength and bond length observed in this table:

Naming Covalent Molecules DIATOMIC MOLECULES – Gaseous elements

HOBrFINCl stands for H2, O2, Br2, F2, I2, N2, Cl2

ALL are diatomic molecules, whose name is the same as the element name

Eg chlorine Cl2

10. Molecular Elements Some elements combine with themselves to form

molecules and this is the way that they are found in nature.

H2(g) hydrogen gas P4(s) phosphorus F2(g) fluorine gas

N2(g) nitrogen gas S8(s) sulfurCl2(g) chlorine gas

O2(g) oxygen gasBr2(l) bromine

I2(s) iodine “HOBrFINCl”

Binary Acids and their Gases There are 5 common binary acids, that form

when the corresponding gas dissolves in water to form an acidic solution. The gas is given a different name from the acid so that we can tell them apart and we use subscripts indicating their state with the formula to tell them apart.

The binary acids consist of HYDROGEN and a second NON-METAL. Their acids end with the suffix -ic.

Hydrogen Halides and their corresponding Binary Acids

Hydrogen Halide Gases (g)

Binary Acids (aq)

HF(g) Hydrogen fluoride

HF(aq) Hydrofluoric acid

HCl(g) Hydrogen chloride

HCl(aq) Hydrochloric acid

HBr(g) Hydrogen bromide

HBr(aq) Hydrobromic acid

HI(g) Hydrogen iodide

HI(aq) Hydroiodic acid

H2S(g) Dihydrogen sulfide

H2S(aq) Hydrosulfuric acid

General Rules for Naming Covalent Molecules

Write the LEAST ELECTRONEGATIVE element first in the name and in the formula.

Naming Method #1: Common method uses the Prefix system ( also called Classical) the

1=mono 2=di 3=tri 4=tetra5=penta 6= hexa 7= hepta

e.g. CCl4 carbon tetrachloride

IUPAC and covalent molecules Naming Method #2: IUPAC system uses

Roman Numeral to indicate the oxidation number of first element. Refer to text’s periodic table for possible oxidation numbers of nonmetals.

CCl4

carbon (IV) chloride

Electronegativity Linus Pauling developed the concept of

electronegativity (En) It is a measure of how strongly an atom

attracts electrons to itself in a covalent bond

Fluorine has the highest En value, and Pauling assigned it an arbitrary value of 4.1

Elements to the left and below fluorine have decreasing En values

Electronegativity TrendsElectronegativity Trends

Ionic or Covalent? Find ΔEn In bond formation, it is useful to look at the

electronegativity difference (ΔEn)

When Pauling looked at a range of bonds and their ΔEn values, a pattern was noticed

Bonds with an ΔEn greater than 1.7 tended to exhibit ionic characteristics

Bonds with an ΔEn below 1.7 tended to exhibit covalent characteristics

HBr En (hydrogen) = 2.1 LiF En (lithium) = 1.0

En (bromine) = 2.8 En (fluorine) = 4.1

ΔEn = 2.8 – 2.1 = 0.7 ΔEn = 4.1 – 1.0 = 3.1

HBr is polar covalent LiF is ionic

Electronegativity differences between 0.4 and 1.7 mean the ELECTRON PAIR WILL NOT BE SHARED EQUALLY, BUT WILL BE closer to the nuclei of the more electronegative element. This creates a POLAR COVALENT BOND.

The O-H bonds in water are polar covalent

There are two types of covalent bonds Atoms that have the same En will have an

ΔEn of zero. These atoms will attract the shared

electrons equally, and so the distribution of electrons is uniform

These are nonpolar covalent bonds

Covalent bonds that have two different elements will have different En values and so the electron distribution will be non-uniform

These bonds are called polar covalent, since one end of the bond will be slightly electronegative (δ-)since the electrons are attracted more to the atom at that end

Valence Shell Electron Pair Repulsion theory (VSEPR) allows us to predict the 3-D shape of a molecule

VSEPR theory states that bond pairs of electrons repel one another, and lone pairs of electrons take up more space than bond pairs

There are four basic shapes which are common in organic molecules:

Linear Bent or V-shaped Tetrahedral Pyramidal

Linear

or

Bent

Pyramidal

Tetrahedral

How do we determine if a molecule is polar or nonpolar?

A polar molecule has an uneven distribution of electrons. This occurs when◦ There is at least one polar bond ◦ The shape of the molecule is asymmetrical◦ Or the shape is symmetrical but the atoms

surrounding a central atom have different En values

Methane: CH4

VSEPR diagram:

Polar bonds? YesOverall dipole? NoMethane is: NON-POLAR

Ammonia: NH3

VSEPR diagram:

Polar bonds? YesOverall dipole? YesAmmonia is: POLAR

Water: H2O

VSEPR diagram:

Polar bonds? YesOverall dipole? YesWater is: POLAR

Carbon dioxide: CO2

VSEPR diagram:

Polar bonds? YesOverall dipole? NoCarbon dioxide is: NON-POLAR

The particle theory states that there are forces between particles, and the forces increase as the particles get closer.

These are the intermolecular forces Compared to covalent and ionic bonds, they

are very weak – but when there are many, they add up to a significant force

Collectively they are called van der Waals forces, but there are three different forces.

These forces have an effect on the boiling point and the solubility of substances.

vanderWaal’s forces vanderWaal’s forces of attraction occur

when the protons in one atom or molecule attract the electrons in a neighbouring atom or molecule.

Since all particles have protons and electrons, all substances have vanderWaal’s forces

Larger molecules have more protons and electrons, and so have stronger vanderWaal’s forces.

vanderWaal’s forces of nonpolar hydrocarbons affect boiling points:

When comparing the boiling points of hydrocarbons (non-polar molecules), we see that the boiling point increases as the number of carbons increases.

Why is this?

Occurs only in polar molecules that have hydrogen and at least one of the following atoms: N, O or F.

These highly electronegative atoms have lone pairs of electrons which are attracted to the hydrogen atoms in neighbouring molecules.

These hydrogen atoms are essentially a proton

Polar substances have a slightly electronegative end and a slightly electropositive end.

Dipole-dipole forces occur when oppositely charged poles momentarily attract one another

Water is not an organic molecule but is essential for life on this planet

All cells are surrounded inside and out with water – anything that interacts with a cell must first be dissolved in water

Physical properties:◦ colourless and transparent◦ liquid at room temperature◦ density = 1.0 g/mL◦ m.p. = 0℃ b.p = 100℃

water has LD, D-D forces, and H-bonding

Water has cohesive properties – the high number of intermolecular forces causes water molecules to ‘stick’ together

Examples:◦ surface tension – beading of water◦ water striders – too light to break surface tension◦ transpiration in plants – transport in xylem tubes

Water has adhesive properties – it’s polar nature causes it to stick to other substances

Examples:◦ capillary action – water ‘climbs’ up small diameter

tubes, or ‘bleeds’ through the microscopic pores and channels in paper or other porous substances

◦ this is due to the hydrogen bonding interactions between the water and the surface of the tube (either SiO2 or the cellulose tubes of paper)

◦ This helps to explain the meniscus inside a tube

Water has outstanding solvent properties Used to be called the ‘universal solvent’, but this

is not a good name, since not everything dissolves in water

The polar nature of water allows any other polar substance or any charged particle to dissolve easily

The δ- will attract the δ+ end of solutes, and this attraction will remain once the solute is dissolved.

The same is true for ionic substances – the cation will be attracted to the δ- end of water, and the anion will be attracted to the δ+ of water.

Water’s density decreases as it changes from liquid to solid.

This is because the distance between molecules in a crystal lattice (as ice) on average further than when in a liquid.

Special Properties of waterSpecial Properties of water