As chemistry unit one point 3

oppositely charged ions in a lattice

Shared pair of electrons

Dative covalency

involves a lattice of positive ions surrounded by delocalised electrons 

the power of an atom to withdraw electRon density from a covalent Bond 

symmetrical

different extents 

permanent dipole-dipole , induced dipole-dipole forces and hydrogen bonding 

boiling points of compounds and the structure of some solids such as ice 

energy

ionic , metallic , giant covalent (macromolecular) and molecular 

• sodium chloride
• magnesium
• diamond
• graphite
• iodine
• ice  

the type of structure and bonding present 

charge clouds 

molecules and ions , limited to 2 , 3 , 4, 5 and 6 coordination 

lone/pair boding repulsion , which is greater than bonding pair/boding pair repulsion 

forces of attraction 

when atoms lose or gain electrons which which is when a metal and non metal react to form a compound

• lose electrons forming cations
• for example a lithium atom loses 1 electron to form a lithium ion 
• Li ---> Li⁺ + e^- 
• the lithium ion has electronic configuration 1s²

isolectronic (has the same electronic configuration) with helium

• gain electrons forming anions . 
• for example a fluorine atom gains an electron to form a fluoride ion . 
• F + e^- ---> F^- 

1s²2s²2p⁶ 

isoelectronic with neon

the electrostatic attraction between cations and anions

lattice

lattice structures

ions of opposite charge , this structure is repeated throughout the ionic compound , as a result ionic compounds are said to have giant structures

molecules

molecular formula

the ratio of cations to anions present in the ionic lattice

the formulae of the cations and anions present in the ionic lattice . in most cases you work out the charges of cation and anion using the periodic table

same as the group number of the metal

the group number of the atom minus 8

H⁺

Ag⁺

• NH₄⁺

Cu²⁺  or Cu³⁺

Zn²⁺

Pb²⁺

Fe²⁺ or Fe³⁺

OH^-

NO₃^-

O^2-

S^2-

SO₄^2-

CO₃^2-

PO₄^3-

CrO₄^2-

neutral

• Al³⁺ ions and SO₄ ^2- ions 
• Al₂(SO₄)₃

ions are fixed in position by strong ionic bonds so are not able to move and carry charge however when molten or dissolved in water ions are free to move

giant lattice structure held together by strong forces of electrostatic attraction between oppositely charged ions which requires a lot of energy to overcome

small displacement causes contact between ions with the same charge , the ions therefore repel and the structure shatters

polar water molecules pull ions away from lattice and cause it to dissolve

• LiCl because both compounds have a chloride ion 
• however Li⁺ is a smaller ion (same charge) and so has a higher charge density . This means that there is a greater force of attraction between cations and anions meaning LiCl requires more energy to overcome its forces of attraction 

Both have O^2- ions however MgO has Mg²⁺ ions whereas Na₂O has Na⁺ ions

Mg²⁺ evidently has a higher charge and is also a smaller ion and has a higher charge density .  This means that there is a greater force of attraction between cations and anions meaning MgO requires more energy to overcome its forces of attraction

metallic bonds 

the electrostatic force of attraction between metal ions and the sea delocalised electrons in a metallic lattice . the delocalised electrons come from the highest energy level

• in a metallic bond 
• each metal ion forms a positive ion
• the positive ions are arranged in a lattice structure 
• the ions in the structure are very close together so the electrons that are lost when the metal forms ions are delocalised 
• delocalised electrons aren't attracted to any particular ion 
• as a reasult these electrons are free to move when a voltage is applied . Pushing electrons in one end of the piece of metal causes electrons to come out of the other end

• regular arrangement of ions of opposite charge
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• regular arrangement of cations surrounded by a sea of delocalised electrons
• 

• the charge (density) of the ions in the lattice -the larger the charge (density) the stronger the bond 
• the number of electrons in the sea of delocalised electrons - the more electrons there are the stronger the bond

• Increases across a period as more electrons become delocalized 
• Decreases down a group as the atomic radius increases - (the charge density decreases)

the attraction between the Mg²⁺ ions in the lattice and the sea of delcoalised electrons is greater than between the Na⁺ ions in the lattice and the sea of delocalised electrons . this is because Mg²⁺ has a higher charge and more delocalised electrons

because there is a stronger metallic bond in lithium because lithium has a higher charge density so there are stronger forces of electrostatic attraction between the metal ions and sea of delocalised electrons

alumnium has a higher charge density and more delocalised electrons . so there are stronger forces of electrostatic attraction between the metal ions and sea of delocalised electrons in aluminium so it has stronger metallic bonds which require more energy to overcome

in a close packed structure

hexagonal close packed (HCP) and simple cubic

stacked directly on top of each other

efficient with 74% of the volume fixed with particles

• the ions in the first layer are arranged so that they are touching each other
• in the second layer the ions sit in the hollows between the ions in the first layer 
• in the third layer the ions are in hollows in the second layer , directly over ions in the first layer

space filling model or unit cell

the metal ions in the lattice are very close to each other . heating one end of a piece of metal makes the ions at that end vibrate more , and these vibrations are passed along the piece of metal 

metals have a giant structure and there are strong forces of electrostatic attraction between the cations and sea of delocalised electrons 

No bonds holding ions together and ions can slide over each other if a large enough force is applied . the strength of metallic bond stops the attraction being broken completely 

because of their metallic bonding , there are strong forces of electrostatic attraction between cations and the sea of delocalised electrons 

strength of metallic bonds

accepting electrons from metal atoms or sharing pairs of electrons 

identical atoms or different atoms 

outer energy level electrons 

pair of electrons is shared between two atoms 

the attraction between each nucleus involved in the bond and the pair of electrons 

dot and cross diagrams . they can also be shown as a straight line between the atoms , this is called displayed formula 

sharing more than one pair of electrons . these are represented in molecular formulae by multiple lines between atoms . double bonds are stronger than single bonds and triple bonds are stronger than single bonds 

a shared pair of electrons

two shared pair of electrons 

three shared pairs of electrons 

• oxygen 
• carbon monoxide 

nitrogen 

lone pairs of electrons 

• they are important in determining the shape of molecules 
• they can also influence the chemical properties of molecules 

• nitrogen is in group 5 and form 3 single bonds with hydrogen 
• 

• since oxygen is in group 6 and forms two covalent bonds with hydrogen 
• 

• co-ordinate 
• dative 

shared pair of electrons in which both electrons are contributed by one of the atoms in the bond 

arrow 

atom with an empty orbital and an atom with a lone pair of electrons . these bonds can only occur in simple molecular substances 

• the coordinate bond is is between the nitrogen atom and one of the hydrogen atoms . it has formed because nitrogen has a lone pair of electrons and hydrogen is able to gain a vacant orbital 
• 
• 

• a coordinate bond forms because oxygen has a lone pair of electrons and hydrogen has a vacant orbital  
• 
• 

• 
• a coordinate bond forms because oxygen has a lone pair of electrons and carbon has a vacant orbital  

simple molecules . these are discrete molecules made from a small number of atoms . 

there are no ions or delocalised electrons involved 

there are weak intermolecular forces 

• strong covalent bonds
• Intermolecular forces between the I₂ molecules

simple covalent (molecular) or giant covalent structures (macromolecular) note we have just been looking at simple covalent

different structures of the same element 

4

bond to itself 

diamond and graphite both of which are allotropes of carbon 

arrangement of atom is repeated many times 

• each carbon atom forms 4 covalent bonds 
• the shape around each carbon atom is tetrahedral 
• the C-C bond angle about each atom is 109.5^O
• the tetrahedral structure is repeated around each carbon atom making diamond extremely hard 
• 

• each carbon atom forms three covalent bonds 
• the shape around each atom is trigonal planar 
• the C-C bond angle about each atom is 120^o
• the carbon atoms form a flat lattice structure made from hexagons 
• the extra electron from each carbon is contributed to a delocalised sea of electrons between the layers 
• 

van der waals forces exist between the layers . these weak forces allow the layers to slide over each other 

it's strong because there are strong covalent bonds between the carbon atoms in each layer and its lightweight because the layers are quite far apart which gives it a low density 

• Graphite has 1 delocalized electron per carbon atom as it only forms three bonds so can
• conduct electricity along the hexagonal sheets. 

because there are strong covalent bonds in hexagonal sheets , this requires a large amount of energy to overcome 

there are covalent bonds which are difficult to break 

strong covalent bonds in a crystal lattice structure  

rigid crystal structure 

no delocalised electrons or ions 

covalent bonds are too difficult to break 

sharing of one or more pairs of electrons 

the nuclei of the two atoms involved in the bond and the pairs of electrons  

equal e.g. H₂ , Cl₂

• unequal e.g. H-F 

fluorine has more protons 

the ability of an atom to withdraw electron density from a covalent bond 

the more electronegative element has a greater share of electrons 

at the tops of groups 5,6,7 

• electro-negativity increases across the periodic table as the number of protons in the nucleus increases . this increases the ability of an atom to attract electrons
• electro-negativity decreases down the periodic table as the amount of electrons in complete energy levels increases . As a result the nucleus is shielded and has less ability to attract electrons 
• fluorine is the most electronegative element owing to its small size  

a pair of electrons is not shared equally 

a covalent bond between atoms with different electro-negativities 

δ^-

δ⁺

• the chlorine atom is more electronegative so has a partial negative charge 
• the hydrogen atom therefore has a partial positive charge 
• as a result the H-Cl molecule can be described as polar H^δ+ - Cl^δ-

• the oxygen atom is more electronegative so has a partial negative charge 
• the carbon atom has two partial charges positive charges as it is bonded to two oxygen atom 
• as a result , the C=O bond can be described as polar  O^δ-=C^2δ+=O ^δ-

• draw the molecule (in 3D if necessary remembering about the influence of lone pairs)
• label any polar bonds using the δ+ , δ- convention 
• then examine the shape of the molecule 
• if molecule has a positive end and a negative end then the molecule is polar but if it doesn't the molecule is non polar 

opposite charges separated by a short distance in a molecule or ion 

the bigger the dipole (difference in charge between atoms) and the more polar the bond 

having ionic character 

ionic rather than covalent 

• large difference in electro-negativity
• electrons completely transferred
• electrons not shared   

• no difference in electro-negativity
• electrons not transferred
• electrons shared equally 

there is a larger difference in electro-negativity 

polarity 

• small highly charged positive ion (these have high charge density) 
• the positive ion is said to be polarizing 
• the negative ion is said to be polarized 
• if this happens to a large enough extent the ionic bond takes on a covalent character 

low charge density which means they are not able to polarise other ions that easily so its compounds have an ionic character 

a high charge density so they are able to polarise large negative ions so many of its compounds have a covalent character 

weak attractive force between molecules . 

the melting points and boiling points of substances and can influence some of their other properties to 

substances together in solid or liquid form .

attractive force that exists between polar molecules 

a molecule in which the charge isn't symmetrically distributed so that one area is slightly positively charged and another negatively charged 

different electron density within them . these molecules are described as polar . so permanent dipole dipole forces are intermolecular forces (forces between molecules) in polar molecules . they only exist between any two molecules that have permanent dipoles.  

• the electronegative chlorine of one HCl molecule will attract the electro positive hydrogen of another HCl molecule 
• the doted line shows the permanent dipole dipole force 

• hydrogen bonding 
• van der waals 

• hydrogen bonding 
• permanent dipole - dipole forces
• van der waals forces 

especially strong permanent dipole-dipole force which exists between very electronegative elements 

• intermolecular force between a lone pair of electrons of an N, O , or F atom in one molecule , and a H atom joined to an N , O or F atom in another molecule 

the dipole-dipole force of hydrogen bonding between the molecules is especially strong 

• water 
• ammonia
• hydrogen fluoride 

• labelled dipoles on each molecule 
• lone pairs on the , O , N or F molecule 
• a dotted line to represent the hydrogen bond 

• with a lone pair of e- on every O
• 

• with a lone pair of e- on every F 
• 

ice has a very regular structure held together by hydrogen bonding between molecules 

• ice floats on water . It is the only substance for which this is the case 
• water has a large surface tension , caused by a network of hydrogen bonds on the surface 
• water has a higher boiling point than would be expected . this is because hydrogen bonds must be broken when water is boiled 

force of attraction between a temporary dipole one one molecule and an induced dipole on another molecule 

non polar and polar molecules but are the dominant force between non-polar molecules 

the asymmetrical distribution of the electron pair in a covalent bond  

constantly moving 

• the electron cloud around an atom or within a non polar molecule is not static 
• at any instant in time the distribution of the electrons may be uneven although on average they are distributed evenly . 
• as a result a non polar molecule may have a temporary dipole 

a dipole to form in a nearby atom or molecule . this dipole is called an induced dipole . The induced dipole can then induce a dipole in a neighbouring atom or molecule . The net effect of this is a force of attraction between the particles called a temporary dipole-induced dipole force 

van der waals forces 

• non polar atoms or molecules 
• diatomic molecules (two atoms joined together with a covalent bond)
• solid crystalline 
• regular arrangement 
• between molecules 
• weak so little energy is needed to overcome them 
• iodine has a low melting and boiling point and can sublime 

it can change directly from a solid to a gas 

• the size of the atom or molecule 
• the area of contact between the atoms or molecules  

• the size of the halogen molecules increases down the group
• the number of electrons in each molecule increases down the group
• as a result temporary dipoles form more readily in the large halogen molecules
• in addition dipoles are more readily induced in adjacent molecules
• as a result the van der waals forces get stronger down the group .
• as a result more energy is needed to overcome these forces  
• the same trend occurs in the alkanes as the number of carbon atoms in the chain increases 

chain branching . this reduces the area of contact between other molecules and reduces the strength of the van der waals forces 

changed state 

changes in energy 

break the forces between the atoms molecules or ions involved 

the energy provided to the system is to break the force 

the attraction between the lattice of positive charge ions and the de localised electrons . this is a strong force so requires a huge amount of energy , which means metals have a high melting point . the melting points increase as the charge on the metal ion and the number of electrons increases 

ions of opposite charge . this attraction is strong so ionic substances are hard to melt 

intermolecular force between the molecules . there are three types of force that can be broken . each of these is weaker than the covalent bond that exists between the atoms within the molecules 

covalent bonds . melting these substances involves a large amount of energy many strong covalent bonds must be broken in order to change state 

• both C=O and H-C are polar bonds since electro negativity N>H and O>C 
• hydrogen bonding between H and =O or N in different molecules 
• using a lone pair of e- on an O or N atom

Sodium chloride