Chapter 3 - bonding and periodicity

Chemical bonds - force of attraction, metallic bonds in metals, covalent bonds in non-metals, generally ionic bonds in non metal and metal.

Metal atoms lose and non-metal atoms gain electrons to obtain a stable electron configuration, often that of a nobel gas, becoming ions. Some d block metals form ions with a nobel gas configuration but most don't.

Formation of ion often shown by dot and cross diagram. Usually only electrons in highest energy shell shown.

Ionic compounds - compounds made up of ios, with strong electrostatic forces of attraction between ions, ionic bonds. Ions in ionic compund form regular 3-D structure, a lattice. Each ion surrounded by ions of opposite charge, repeated many times forming giant structure.

Ionic compounds don't exist as molecules, compound's molecular formula thererfore an ionic formula. When working ou the formula remember you nedd equal numbers of positive and negative charges.

Metallic bonding occurs in metals, with atoms closely packed in solid state. Highest occupied energy levels overlap causing electrons in outer shell to become delocalised - free to move through structure.
Atoms then effectively positively ionized to form cations whcih are regularly arranged to form lattice. Positive ions don't repel as strong electrostatic forces of attraction between ions and delocalised electrons, metallic bonds.

Metals are malleable (hammered into shape) and ductile (stretched to make a wire) as layers of cations can slide over each other if sufficient force applied but metallic bonds generally stop stucture being destroyed.

Properties of metals changed by adding other metals - alloy. Some alloys much togher than individual metals they contains but some alloys also more easily shaped than pure metal.

Magnesium, zinc and many other metals have a hexagonal close-packed structure. one of most efficient ways to arrange particles with 74.1% of space filled with particles. Two alternating layers in strucutre -fills gaps made from previous layer.

Covalent bond - shared pair of electrons from highest occupied energy level of atoms involved. Both nuceli attarcted to electron pair, with lots of energy needed to overcome this electrostatic force of attraction -> strong. Substances with covalent bonds form molecules.

Simple molecules - msall number of atoms joined by covalent bonds, hydrogen smallest molecule like this, shares its 1s1 electron to have electron configuration 1s2 becoming isoelectric with helium.
Can use dot and cross diagram to show bonding or straight line to show covalent bond - displayed or graphical formula.

Covalent compounds - covalent bonds between atoms of diff elements, shown by dot & cross diagrams/display formulae. Elecs in highest energy level not involved in bonding - lone pair.

Diamnd and graphite - allotropes (same state but different structures) of carbon. Carbon gain 4 elec, so forms 4 covalent bonds. diamond each atoms joined to 4 others, bonds forming tetrahedral arrangement. Reapeated many times so diamond is giant covalent/macromolecular structure. Diamond - hardest natural substance due to strucutre and bonding.

Co-ordinate/dative covalent bond - one atom contributes both shared electrons with the other atom contributing nothing. Once formed behaves in the same way as a covalent bond. For a dative bond to form there must be: a lone pair of electros on 1 atom; a vaccant orbital on the other. When drwaing display formula of co-ordinate bond, line replaced by arrow pointing towards atom with vacant orbital.

Aluminium chloride - covalent compound. Al has 3 electrons in outer shell, each Cl in AlCl3 has 3 lone pair of electrons, Al atom has vacant orbital so Cl atom in 1 molecule of AlCl3 contributes elect to form dative bond with Al in 2nd AlCl3 molecule, Cl atom from this molecule forme dative bond with 1st AlCl3 molecule forming a dimer with 2 dative bonds.

Polyatomic ion - charge particle consisting of 2 or more atoms joined by covalent bonds, e.g. ammonium ion, NH4+ formed from ammonia and H+. Contains 4 covalent bonds, one being a dative bond. Ammonia has love pair of electrons on N atom and H+ has vacant orbital as lost only electron. so N atom contributes lone pair.

Work out shape of molecule by counting electron pairs, each one acting as negative cloud repelling each other to be as far away as possible, keeping force of repulsion to minimum, givig each molecule a characteristic shape.

Linear molecule - atoms in straight line, bond angle of 180 (all diatomic molecules are linear).

Trigonal Planar molecules - 3 bonding pairs of elec in trigonal planar shape, bond angle 120.

Tetrahedral molecules - 4 bonding pairs of elecs in tetrahedral shape, bond angle109.5.

Trigonal bipyramidal - 5 bonding pairs of elec in trigonal bipyrimidal shape, 2 diff bond angles. has two trigonal pyramidal shapes around equatorial atom, lie on plane, bond angle 120, atoms at 2 points, axial atoms and equatorial atoms, bond angle 90.

Octahedral molecule - 6 bonding pairs of elec in octahedral shape, bond angle 90.

There may be lone pairs of electrons which affect the bond angles. Pairs of electrons repel each other and lone pairs are more compact so have greater repulsion -> reduce bond angle.

Darwing molecules - linear and trigonal planr molecules are flat so easy to draw.

Tetrahedral: ammonia, has 3 bonded and 1 lone pair of electrons. Lone pair repels more strongly so angles reduced to 107, molecule is trigonal pyramidal; water, 2 bonding and 2 lone pairs of electrons so bond angle reduced to 104.5, shape of molecule is bent line.

Trigonal bipyramidal arrangements: sulfur tetraflouride, 4 bonding and 1 lone pair of electrons, reduces bond angle to 118; chlorine triflouride, 3 bonding and 2 lone pairs of electrons bond angle reduced to 90 and 180.

Octaherdral arrangements: sulfur hexafluoride, 6 bonding pairs, bond angles of 90; iodine pentaflouride, 5 bonding and 1 lone pair of electrons, square pyramidal.

General method for working out shapes - central atom bonded to atoms that can only form single bonds, simple method for working out shape:
Step 1: Find group number of central atom & add number of atoms around central atom to it
Step 2: If ion negatively charged, add charges to and if positively charged subtract charges from number in step 1
Step 3: Divide number from step 2 by 2 giving number of electron pairs aound centre atom letting you work out arrangement of electron pairs : 2 pairs -> linear, 3 pairs -> trigonal planar, 4 pairs -> tetrhedral, 5 pairs -> trigonal bipyramidal, 6 pairs -> octahedral.
Step 4: lone apirs later shape and bond angles. Number from step 3 - number of atoms around cnetral atom = number of lone pairs.

Two bonded paris of electrons in a covalent bond different elements one of bonded atoms larger share of electrons if difference in electronegativity.

Electronegativity - tedancy of an atom to pull electrons to its nucleus. Bonded atoms, equal electronegativity bonding electrons shared equally. Different electronegativity, more electronegative element, largershare of electrons.
Electrogenativity cannot be be measured, has to be calculated, giving Pauling electronegativity scale.

Bonding pair of electrons not equally shared, polar bond formed. More electrongeative atom, nagative charge, other positive charge. Molecule then has dipole, pair of saparated charges of opposite signs.

Non-polar molecules: If molecule is symmetrical, no net dipole as overal dipoles are cancelled -> even if molecule has poalr bonds, is non-polar.

Polar molecules: Molecule will be polar if: conatins polar bonds; has a net dipole. Water polar as atoms don't lie in straight line so dipoles not cancelled out.

Electronegativity and bond character - bigger difference in electronegativity, bigger the dipole, more polar the bond. As electronegativity increases, covalent bond apapts covalent character. v.large difference, bonding becomes ionic not covalent.

Ionic bonds show polarity. Electron cloud around negative ion can be distorted and drawn to positive ion -> negative ion polarized by positive ion. Negative ion sufficiently polarized, ionic bond adapts covalent charcter, level depending on: charge and size of negative ion as more likely to be polarized if big; charge density of cation, high charge, small large polarizing power.

Sodium ions large charge density but don't polarize cholride ions so forms ionic bonds. Mg ions, large charge density so polarize Cl ions -> MgCl3 adapt covalent characteristics, so becomes covalent.

Forces bewteen molecules, intermolecular forces, attract molecules to each other, weakest intermolecular force, van der Waals'.

Van der Waals - forces of attraction bewteen a temporary dipole and an induced dipole on polar and non-polar molecules.  Electron cloud around non-polar molecule or atom, not static, sometimes unevenly distributed, -> temporary dipole. Temporary dipole in one molecule can cause dipole in enighbouring molecule -> induced dipole. Opposite charges to neighbours tempoary dipole, causes attractive intermolecular force.

Iodine - diatomic molecules, I2 joined by covalent bond. Molecules regualrly arranged -> molecular crystal. Van der waals between molecules. Iodine can sublime or turn directly from solid to vapour.

Strenght of van der waals depends on: size of atom or molecules; area of contact between molecules.  As atom increases in size, more electrons -> more easily polarizied -> temporary dipoles form easier -> induced dipoles form more easily. As area of contact increases, strength of Van der Waals increases.

Permanent dipole forms when 2 atoms in covalent bond have different electronegativites. Molecules with permanent dipole have dipole-dipole and VDW forces between them. Some molecules, dipoles cancel out -> non-polar, have no permanent dipole-dipole forces or net dipole. Dipoles don't cancel out, net dipole -> dipole-dipole and VDW forces.

Hydrogen bonds - permanent dipole-dipole forces between a hydrogen atom covalently bonded to flourine, nitrogen and oxygen atoms that have a lone pair of electrons. F, O and N, v.electronegative -> form v.polar covalent bond with H, pull H electron away strongly attracting H.

Ice - water forms two hydrogen atoms covalenty bonded to oxgen. Oxygen has 2 lone pairs -> hyrdogen bond forms, forming along same axis as H-O bond on neighbouring water molecule -> ice has regular open lattice structure. Molecules further apart then when in water -> ice has lower density.

Three states of matter: Solids, particles close together in regular arrangement with strong forces, able to vibrate;Liquids, particles close but fewer forces acting on them so move randomly around each other;Gases, particles far apart, no forces on them, move randomly.

Changing state - Melting: solid heated particles gain energy, vibrate increasingly , and temo increases. Melting point, energy enough to break some intermolecular foces, structure breaks down -> liquid. Evaporating: liquid heated, particles gain energy, move increasingly, some particles enough energy to break all bonds -> move away as gas. Boiling: bubbles of vapour form inside liquid at boiling point. Energy enough to break all bonds between particles.

Compairing different substances - Ionic substances: conatin ions attracted by electrostatic forces. Ionic bonds strong and many, lot of energy needed to break. Metallic substances: regular lattice of positive ions attracted to surrounding cloud of electrons by electrostatic forces. Bonds strong and thoroughout structure, lots of energy to break. Giant Covalent substances: atoms covalently bonded, bonds strong and many, lots of energy to break. Molecular substances: simple, molecules attracted by weak forces, little energy needed to break -> melting point low. Strong covalent bonds between atoms, need to be broke to boil. Strong -> more energy needed but still low.