chemistry alevels - electrons, bonding and structure

another name for shells

energy levels

feature of shells

as the shell number increases so does the energy

define principle quantum number

the shell/energy level number

electron shell

a group of atomic orbitals with the same principle quantum number

sub shell

a group of orbitals of the same type within a shell

p - orbital

dumbell shape

1st shell

n and electrons

2

2(n²)

2nd shell

8

3rd shell

18

4th shell

32 electrons

define first ionisation energy

the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1⁺ ion

example of first ionisation energy

• Mg (g) —> Mg⁺(g) +e^-

• Cl (g) —→ Cl⁺ (g) +e^-

the trend in the ionization energy in the first 20 elements in a period

across a period the ionisation energy generally increases, due to the nuclear charge increasing, increasing the number of protons, resulting in a stronger attractions to the electrons on the outershell to the nucleus. therefore more energy is required for ionisation

factors affecting ionisation energy

• atomic radius

• nuclear charge

• electron shielding

atomic radius

the greater the distance between the nucleus and the outer electron - the weaker the attraction - smaller the ionisation energy

the smaller the distance between the nucleus and the outer electron -the stronger the attraction - higher the ionisation energy

nuclear charge

the more number of protons in the nucleus, the greater the attraction between the nucleus and the outer electron - the higher the ionisation energy

electron shielding

repulsion between the inner electrons and the outer electron is shielding - this reduces attraction between nucleus and outermost electron - reduces ionisation energy

s orbital - features

• spherical in shape (not cirular or round)

• all shells contain 1s orbital

• max of 2 electrons (2×2)

p orbital - features

• dumb bell shaped orbital

• from shell 2, each shell has 3x p - orbitals

• max of 6 electrons (3×2)

d orbitals - features

• more complex shapes

• 3rd shell each has 5d orbitals

• max of 10 electrons (5×2)

Hunds Rule

as orbitals have lots of different shapes boxes are used to represent them as easier

how much does these subshells hold s p d

2 6 10 electrons

electronic configuration order

1s 2s 2p 3s 3p 4s 3d 4p 4d

aufbau principle

short way to write the electronic configuration

F: (1s²)2s²2p⁵ so is (He)2s²2p⁵

USE NOBLE GASES AT THE BEGINNING THE NEAREST ONES

electronic configurations of Cl^-

1s²2s²2p⁶3s²3p⁶

how many full orbitals are in atom of sulfur

7 - 2 electrons in each orbital as sulphur has 14 electrons

remember boxes

ionic bond

transfer of electrons between metals and non metals

the metals in an ionic bond

is oxidised so it loses electrons and becomes a positive ion

non metals in an ionic bond

is reduced so it gains electrons so becomes a negative ion

structure of ionic compound

electrostatic force that acts in all directions in an ionic bond

tightly packed structure called crystal

3D lattice

structure of alternating anions and cations

physical properties of ionic compound

• do not melt or boil very easily because the electrostatic forces are very strong. as solids they are hard and brittle

• forms a crystal lattice that results in release of energy, making the resulting salt more stable then before

properties of ionic compounds

when salts are dissolved in water, this breaks the ionic bond and allow the individual ions to move freely in water forming electrolyte

ionic bond energy - lattice energy - the energy required to seperate one mole ( a chemical quantity) of ions is an ionic compound into its gaseous ions.

the lattice energy is affected by

• the size of the ions - smaller the ions in bonds, the greater the lattice

• the charge on the ions - greater the charges in the bonds, the greater the lattice energy

sodium chloride, potassium chloride, magnesium chloride

rank them from the highest to lowest melting points with reasons

1. magnesium chloride - ionic radius is smaller so greater charge density so electrostatic force is stronger so more energy needed to overcome the force

2. sodium chloride

3. potassium chloride - 1 extra shell

charge density

charge ions / volume of ions

why does MgO have a much higher boiling point than NaCl

comparisons of cations

Na⁺ and Mg²⁺ ionic radius of Mg ion is smaller than Na+ ion and has a greater magnitude (+2 +1) Mg²⁺ ion forms a stronger electrostatic forces of attraction. chloride ions has 3 shells/energy, whereas oxide has 2. Cl^- ion has larger ionic radius than oxide ion. Cl^- have a lower magnitude of charge -1 in comparison to -2. MgO forms strongest electrostatic forces of attraction.

if energy taken in to break lattice it is an

endothermic reaction

solubility

two processes - ionic lattice must be broken down and water molecules must attract and surround the ions

covalent bond

covalent bonding the attraction is localised, this means that it acts solely between the shared pair of electrons and nuclei of the two bonding atoms

dative covalent bond is also known as

coordinate bonds

define dative covalent bond

higher electronegativity means that the bond is

polar

define polarity

difference in electronegativity between elements

define dipole

consists of 2 equal but opposite charges or magnetic poles seperated by distance

dipole is often seen in

polar molecules with uneven charge distribution, leading to a positive and negative end

properties of covalent compounds

• low melting ang boiling point - the bonds themselves are strong but the forces between the molecules are weak called the intermoleculare forces, so they are easy to break and disrupt

• poor conductors of electricity - doesnt contain charged ions

• soft and flexible - if compounds crystalline not the case

• non polar covalent compounds dissolves poorly in water - water is polar so the rule for dissolving is like dissolves like —> polar and polar. non polar and non polar

define electronegativity

a measure of the tendency of an atom to attract a bonding pair of electrons in a covalent bond

greater electronegativity means

it attracts electrons towards it

factors affecting electronegativity

atomic charge

distance from the nucleus

electron shielding

what is pauling scale

measures electronegativity of atoms

the electronegativity across a period…

increases as the atomic radius decreases, meaning that there is an increase in nuclear attraction and in nuclear charge

the electronegativity down a group …

decreases, this is because the atomic radius increases - number of energy levels - meaning that there is a less of a nuclear charge

which elements have high electronegativity

group 7s, nitrogen oxygen

when the bonding pairs are unequally shared between to atoms this is due to

the difference in electronegativity (polar molecule bond) this sets up a permanent dipole

define permanent dipole

a dipole in a covalent bond that does not change

in a non polar bond the

bonded electron pair is shared equally between the bonded atoms

bond will be non polar when

the bonded atoms are the same or if they have the same or similar electronegativity

non polar bonds will form

pure covalent bonds

non polar solvent example

hexane

polar bonds are

bonded electron pairs shared unequally between the bonded atoms, contains different electronegativities and different bonded atoms

polar bonds forms

polar covalent bonds

example hydrogen chloride

hydrogen has an electronegativity of 2.1 and chlorine has an electronegativity of 3.0. chloride is more electronegative than hydrogen, this mean it has a greater attraction for the bonded pair of electrons than hydrogen.

water is polar

OH bonds have a permanent dipole these act in different directions but not opposing each other

carbon dioxide is non polar

the CO double bond has a permanent dipole and the 2 dipoles acts in opposite directions cancelling each other out having an overall dipole of 0

shapes of molecules: solid line

bond in the plane of the paper

shapes of molecules: solid wedge

the bond is coming out of the plane of the paper

shapes of molecules: dotted wedge

goes into the plane of the paper

there is in repulsion in this order

bonded pair bonded pair

bonded pair lone pair

lone pair lone pair

tetrahedral

4 bonded pairs

bond angle of 109.5

centre atom with 4 connected

pyramidal

3 bonded pair

1 lone pair

bond angle is 107

1 in the centre with 3 other atoms

non linear

2 bond pairs

2 lone pairs

bond angle of 104.5

1 centre atom with 2 joined

linear

2 bonded pairs

bond angle of 180

trigonal planar

3 bonded pairs

bond angle of 120

octahedral

5 bonded pairs

bond angle of 90

electron repulsion theory

predicts the arrangement of electron pairs around the central atom of unfamiliar molecules and ions

shape of molecules: electron pairs

the greater the number of electron pairs the smaller the bond angle as there is more repulsion

lone pairs repels more strongly than bonded paits

electrons pairs tends to repel each other as far as possible

intermolecular forces

in between

3 types of intermoleculare forces

induced dipole dipole interaction (weakest)

permanent dipole dipole interaction

hydrogen bonding (strongest)

induced dipole dipole interactions