Bonding, Structure, Properties And Energy Changes

Front 1

Proton

Back 1

Positive, mass of 1x10-27g, found in nucleus

Front 2

Neutron

Back 2

Neutral, mass of 1x10-27g, found in nucleus

Front 3

Electron

Back 3

Negative, mass of 1x10-31, orbits nucleus. Electrons are 2000x smaller than protons so do not contribute to total mass.

Front 4

Isotopes

Back 4

Same element, different mass number.
Relative Atomic Mass = (Amount/% x mass) x (same for other isotope/s)
The relative atomic mass refers to the mass concerning isotope abundance.

Front 5

Periodic Information

Back 5

Vertical Group - same electron number in valance shell so react similarly
Row - Same no. of shells (more shells more reactive)

Front 6

Valency

Back 6

Combining power of an element e.g. 2,8,8,2
If non-metal it is the no. of bonds an element can form = no. of unpaired electrons

Front 7

Chemical Bonding

Back 7

An electrostatic force of attraction between two atoms (=ve and -ve charge without current or movement - static). Usually means atoms obtain full valance shells and are stable.

Front 8

Ionic Bonding

Back 8

Electrons are transferred from one atom to another. The atoms become charged ions (cations - +ve and anions -ve). Ionic bonds then form due to attraction between the opposite charges. Held together by attraction. Metal <--> Non-Metal.

Front 9

Ionic Formula

Back 9

Ratio of ions of each type

Front 10

Ionic Formula naming

Back 10

-ide: suggests only two element in a compound (except OH-)
-ate: compound contains O
-ite: Some oxygen (less than -ate)
Prefixes hypo- (lease O) and hyper- (most O)
Bicarbonate: Hydrogen Carbonate
Some metals take on more than one value of charge, shown by roman numerals.

Front 11

Difficult ions to remember...

Back 11

Cyanide CN-
Tin Sn+2
Hypochlorite OCl-
Chlorate ClO3
Chromate CrO4

Front 12

Molecular Formula

Back 12

Between two non-metals. The molecule is the smallest particle of an element/compound capable of stable and independent existence. Bonded with covalent bonds and shared electrons.

Front 13

Monatomic and diatomic

Back 13

Monatomic - nobel gases, one atom per mol.
Diatomic - 2 atoms per mol. e.g. O2, H2, N2

Front 14

Octet Rule Exceptions

Back 14

Elements in row 2 may have less than 8 electrons to be stable
Elements in row 3 may have more than 8

Front 15

Dative bonds

Back 15

Shred electron pair provided by the same atom

Front 16

Shapes of molecules with 4 regions of negative charge...

Back 16

4 regions of negative charge repel to the corners of a tetrahedral, 109 degrees. If all are bonding shape is tetrahedral. If 3 are bonded and one is not shape is trigonal pyramid. Is two are bonded and 2 are not shape is bent and if all but one are non-bonding then the shape is linear. These will all be less than 109 degrees.

Front 17

Shapes of molecules with 3 regions of negative charge...

Back 17

3 regions of negative charge repel as far as possible to the corners of a triangle. If all are bonding shape is trigonal planar, 120 degrees. If 2 are bonded and one is not shape is bent. Is 1 is bonded and 2 are not shape is linear. These will both be less than 120 degrees.

Front 18

Shapes of molecules with 2 regions of negative charge...

Back 18

2 regions of negative charge repel as far as possible along a line. The shape will always be linear, 180 degrees.

Front 19

Bond angles

Back 19

Note: Multiple bonds repel more strongly than single bonds as there is a greater charge conc. Also hold the atom closer together
A molecule, if it includes lone pairs, will have a slightly smaller bond angle than if all pair were bonded as when they are not bonding they are slightly closer to the central atom so push bonded pair closer together, repel slightly more.

Front 20

Bond Polarity of identical atoms

Back 20

Non-polar as both nuclei have the same attraction to the shared electron pair so no charge seperation

Front 21

Polarity of non-identical atoms

Back 21

Can share the electron pair unequally if one atom has a greater electronegativity and stronger attraction to the electrons it will become slightly positive and the other atom slightly negative setting up a charge separation known as a bond dipole. The greater the difference in electronegativity the greater the shift of charge density and the greater the dipole strength.

Front 22

If the electronegativity difference of two atoms is more than 1.6....

Back 22

It can be said that the electron pair spends virtually all of its time around one atom with the greater attraction forming ions and eventually ionic bonds.

Front 23

Shape and polarity

Back 23

If the shape is symmetrical then the bond dipoles can cancel so the overall molecule is non-polar but if it is unsymmetrical then they add up and the molecule is polar. Some parts slightly positive and some slightly negative.

Front 24

Metallic Solid

Back 24

Close Packed array of cations surrounded by a uniformly distributed electron sea. Held together with non-directional electrostatic bonds. Metal atoms achieve stability by losing their electrons which form the mobile cloud preventing the positive cations from springing apart due to repulsion from like charges.

Front 25

Metallic Properties

Back 25

Strength varies, many different melting points
Malleable and Ductile as the attraction between the cations and e- is non-directional so layers of atoms can slide over one another while maintaing strength, no like charges introduced
Conduct as electrons are free to move
Closely packed cations so very dense

Front 26

Ionic Structure

Back 26

The cations and anions are held together in a lattice structure and are held together by electrostatic forces. The arrangement of ions depends on their relative size (often situated on planes)

Front 27

Properties of Ionic Compouds

Back 27

High Bp and Mp due to strong electrostatic forces.
Hard and Brittle as if a particular plane is disturbed so that similar charges are together, repulsive forces are produced and the particles come apart.
Only conducts in solution or as a liquid as this is the only time ions are free to move.

Front 28

Solubility of ionic compounds

Back 28

Dissolves in water and other polar solvents as the +ve H ends of the water are attracted to the anions and the -ve O ends of the water are attracted to the cations. This attraction is enough to overcome the forces between ions and the ions are carried into water

Front 29

Molecular Compounds

Back 29

Held together with weak intermolecular forces but strong covalent bonds.

Front 30

Simple Discrete Molecule Properties

Back 30

Low Mp and Bp as only a small amount of energy is required to overcome weak intermolecular forces.
Soft as weak intermolecular forces mean layers slide over one another.
Do not conduct as there are no free electrons.
Polar dissolves in polar solvents and non-polar in non-polar solvents as the forces are especially weak so intermingle easily.

Front 31

Diamond (3-D network MACRO)

Back 31

A network of C atoms joined by strong covalent bonds in a tetrahedral arrangement so have a high Mp and Bp and are hard and brittle. No mobile electrons to conduct as all electrons involved in bonding each C to 4 others. Insoluble in all solvents as there is nothing strong enough to overcome the strong covalent bonds.

Front 32

Graphite (2-D)

Back 32

Each C bonded to 3 others in a layered structure with strong covalent bonds within layer and weak intermolecular between. Good lubricant as the layers can slide over one another. Conducts as there is one e- free to move between layers.

Front 33

Buckminster Fullerene

Back 33

C60 discrete molecular. Each C atom surrounded by 3 others with 1 free electron to conduct, but partially trapped in ball . Non polar so insoluble and hard.

Front 34

Exothermic

Back 34

Making bonds. Heat energy released so surroundings get warmer. Products have less energy than the reactants. The enthalpy change is negative. If exothermic one way will be endothermic the other.

Front 35

Endothermic

Back 35

Breaking bonds. Heat energy is absorbed so the surroundings cool down. The energy to break the bonds is larger then the energy produced making the new bonds. The enthalpy is positive. If endothermic one way will be exothermic the other.

Front 36

The enthalpy change

Back 36

Enthalpy change = enthalpy change of bond broken - enthalpy change of bonds formed/made.
Measures in kJ/mol
If 1 mol is burnt this much energy is required.

Front 37

n = m/M

Back 37

Number of mols (mols) = Mass (g) / Molar mass
3 s.f.