Atoms, Elements and Compounds

C

Co

Cu

Cp

Helium

Neon

Argon

Krypton

It has no charged particles.

The number of protons equals the number of neutrons.

The number of protons equals the number of electrons.

The neutrons cancel out the protons.

Pale yellow

Greenish-yellow

Red-brown

Purple

Ionic bonding

Covalent bonding

Metallic bonding

Hydrogen bonding

7

13

15

17

2, 8, 1

2, 8, 7

2, 8, 8

2, 8, 8, 2

The elements change from non-metals to metals.

The oxides of the elements change from acidic to basic.

The number of outer-shell electrons increases from 1 to 8.

The proton number decreases.

They all have eight electrons in their outer shell.

They are all diatomic gases.

Helium is used to provide an inert atmosphere in welding.

Argon is used in light bulbs because it is non-flammable and unreactive.

Graphite

Silicon(IV) oxide

Brass

Potassium chloride

Identical mass number

Identical number of neutrons

Identical atomic number

Identical radioactive stability

$^{24}_{12}Mg$

$^{12}_{6}C$

$^{24}_{11}Na$

$^{12}_{12}X$

The number of neutrons

The mass number

The number of protons

The density

20

22

42

2

6 protons, 8 neutrons, 6 electrons

8 protons, 6 neutrons, 8 electrons

6 protons, 6 neutrons, 8 electrons

14 protons, 6 neutrons, 14 electrons

6%

94%

50%

75%

It has more electrons.

It has a larger nucleus which affects bonding.

It has a greater atomic mass, leading to slightly stronger London dispersion forces.

It has a different electronic configuration.

17

18

20

37

The mass of an electron is significant.

The presence of various isotopes with different abundances.

The mass of a proton is exactly equal to a neutron.

Errors in measuring the mass of the nucleus.

Isotopes of the same element.

Atoms with the same mass number (isobars).

The same element in different oxidation states.

Allotropes of the same element.

$AgCl$

$Ag_2Cl$

$AgCl_2$

$SiCl$

Group I

Group VII

Group VIII (Noble gases)

Group II

1

2

3

4

Sodium Chloride

Potassium Iodide

Calcium Fluoride

Magnesium Oxide

$H_2O$

$CO_2$

$CaO$

$NH_3$

$1-$

$2-$

$1+$

$2+$

$NaCl$

$MgO$

$AlN$

$CaS$

$Cl^-$

$F^-$

$Na^+$

$O^{2-}$

It does not contain any ions

The ions are held in fixed positions within the crystal lattice and cannot move

The electrons are used up in the ionic bonds

Sodium chloride is a covalent compound in the solid state

Cesium and Fluorine

Lithium and Iodine

Sodium and Chlorine

Potassium and Bromine

They have too many protons

They already have a full outer shell of electrons

They are too small

They are metals

6

8

10

12

They both conduct electricity

They both have very high melting points

They are both very soft

They both have a simple molecular structure

2

4

6

8

$H_{2}O$

$SiO_{2}$

$C_{6}H_{12}O_{6}$

$CO_{2}$

2

18

20

26

Bonding involves sharing electrons.

Does not conduct electricity.

Has a very high melting point.

Contains non-metal atoms.

$CH_{4}$

$NH_{3}$

$H_{2}O$

$BF_{3}$

Linear

Trigonal planar

Tetrahedral

Octahedral

Diamond has ionic bonds, graphite has covalent bonds.

Diamond atoms are larger than graphite atoms.

Diamond has a 3D tetrahedral lattice of strong covalent bonds, while graphite has weak forces between layers.

Graphite contains delocalized electrons which weaken the structure.

They stop moving

They flow toward the positive terminal

They flow toward the negative terminal

They turn into positive ions

Potassium (Group 1)

Sodium (Group 1)

Magnesium (Group 2)

Cesium (Group 1)

Anions

Cations

Isotopes

Allotropes

Always weaker than covalent bonding

Strong attraction between cations and delocalized electrons

Weak attraction between anions and cations

Only exists at very low temperatures

Stronger metallic bonds due to more delocalized electrons and smaller ionic radius.

Weaker metallic bonds due to having more protons.

The same bond strength because they are in the same period.

No metallic bonds because Group 2 elements are non-metals.

High density

Electrical conductivity

Malleability

Lustre

The electrons only move in one direction.

The attraction between the ion and the electron sea acts in all directions.

The metal can only conduct electricity in one direction.

The bonds only form between atoms of the same element.

To break the strong covalent bonds between the metal atoms.

To provide enough energy to overcome the strong electrostatic attractions in the metallic lattice.

To turn the metal into a gas so it can react with carbon.

To remove the delocalized electrons so the metal becomes ionic.

Hardness decreases because more electrons cause more repulsion between atoms.

Hardness increases because more delocalized electrons and higher cation charges result in stronger metallic bonds.

Hardness remains the same because the atomic size increases at the same rate.

Hardness increases because the metals become more non-metallic across the period.

Alloys are mixtures, not compounds.

Alloys always have a higher melting point than the pure metals they are made from.

Alloys can be designed to be more resistant to corrosion.

The conductivity of an alloy is usually lower than that of the pure solvent metal.

1:1

1:2

1:3

1:4

They are arranged in a random way.

They are arranged in layers of hexagons.

They are arranged in a cubic shape.

They are not bonded to each other.

Ionic

Covalent

Metallic

Dative

It is a gas.

The number of atoms depends on the size of the crystal.

It is an ionic compound.

It does not have any bonds.

Only the weak intermolecular forces need to be overcome.

The entire structure must be broken down by breaking strong covalent bonds.

Melting occurs at low temperatures because the atoms are light.

The bonds are broken by the movement of delocalized electrons.

None, it only connects to silicon atoms.

Two, forming an O-O-O chain.

Four, in a tetrahedral arrangement.

One, forming a double bond.

Strong covalent bonds

Ionic attractions

Weak intermolecular forces

Metallic bonds

Graphite has $0.05L$ more covalent bonds than diamond because it contains delocalized electrons.

Diamond has $0.05L$ more covalent bonds than graphite.

Both samples contain exactly $0.10L$ covalent bonds because they contain the same number of atoms.

Diamond has $0.10L$ more covalent bonds than graphite.

Both contain only $C-C$ bonds.

Diamond contains $C-C$ bonds, while $SiO_{2}$ contains $Si-O$ bonds.

Diamond contains $C-C$ bonds, while $SiO_{2}$ contains $Si=O$ double bonds.

Both structures are held together by ionic attractions between atoms.

Graphite

Diamond

Silicon dioxide

All have approximately the same atomic density.