Chemistry- Chemical bonding

define electronegativity

the power of an atom to attract electrons to itself

list of factors that effect electronegativity of an element

1) atomic radius, increase will = decrease in electronegativity

2) shielding, filled energy levels shielding effect of nuclear charge = decrease in electronegativity

3) nuclear charge, increase in positive charge of protons will mean electrons more attracted to nucleus so = increase in electronegativity

trend of electronegativity across and down periodic table

across period: increases across a period due to increasing nuclear charge, shielding becomes negligible

down group: decreases down group as atomic radius is increasing and so is the effect of shielding

define a dipole moment

light charges on an atom in a covalent bond due to differences in electronegativity

define ionic bonding

the electrostatic attraction between oppositely charged ions

what are the structures of ionic compounds

1) giant ionic lattice 2) crystalline solids

what are the properties of ionic compounds

high melting and boiling points

display the ionic bonding for NaCl, MgO, CaF

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define metallic bonding

strong electrostatic forces of attraction between metal cations and delocalised mobile electrons

structure of metallic compounds

is a metallic lattice of positive ions surrounded by mobile electrons

define covalent bonding

the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons

how are the molecules in a covalent compound held together

with weak intermolecular forces

what are the properties of a covalent compound

low melting and boiling points

what can elements in period 3 do with their octets and what does this mean and why

they can expand their octets, this means holding more than 8 electrons on their outer shell, this is due to the vacant 3d orbitals available for bonding

describe (draw) the covalent bonding in: hydrogen, H2 • oxygen, O2 • nitrogen, N2 • chlorine, Cl 2 • hydrogen chloride, HCl • carbon dioxide, CO2 • ammonia, NH3 • methane, CH4 • ethane, C2H6 • ethene, C2H4

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define a dative (coordinative) bond

a covalent bond in which atoms share electrons from the same atom

draw the dative bonding in Al2Cl6

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draw the dative bonding in NH4

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define bond energy

the energy required to break one mole of a particular covalent bond in a gaseous state

how would you be able to determine reactivity of a molecule by looking at its bond energy and bond length

if it has a long bond length= will be easier to break as there is less electron density, will be more reactive

if it has short bond length= will be harder to break as there is higher electron density, will be less reactive

do short bonds have high electron density

yes

what does bond length increase with

increases with atomic radius

what must happen for a covalent bond to form (in terms of orbital overlap)

unpaired valence electrons must overlap, forming a sigma bond

how is sigma bond formed

from the direct overlap of orbitals between the bonding of atoms

what must happen for pi bonds to form

formed from the sideways overlap of adjacent p orbitals above and below the pi bond

define hybridisation

mix of atomic orbitals like s and p to form new hybrid orbitals such as sp, sp2, sp3

sp orbitals: what % is s and p properties and angle do the molecules have that they form

50 % s + p, form molecules with 180 degrees

sp2 orbitals: what % has s and p properties, what degree angle will the molecule have that it forms

33% s property 66% p property, forms molecules with 120 degrees

sp3 orbitals: what percentage are s and p properties, what angle in molecule is formed

25% s, 75% p orbitals, forms molecules with 109.5 degree angles

what bond is sp hybridisation

triple

what bond is formed from sp2 hybridisation

double

what bond is formed from sp3 hybridisation

single bond

what pi and sigma bonds are formed in these molecules: H2, C2H6, C2H4, HCN, N2

H2: 1 sigma bond, C2H6: all sigma bonds, C2H4: 1 sigma 1 pi, HCN: 2 sigma, 1 pi, N2: 1 sigma 1 pi

give a summary of the VSEPR theory

about valence shell electron repulsion where electrons will repulse each other and spread out as far as possible which influences the shape of the molecule

linear: draw shape, bond angle, e density

bond angle = 180 2 areas of e density

trigonal planar

drawn correctly, 3 areas of density, bond = 109.5

tetrahedral

4 areas of e density, angle = 109.5, drawn correctly

pyramidal

4 areas of e density, bond = 107, drawn correctly

non-linear

3-4 e density, bond = 104.5, drawn correctly

octrahedral

6 areas of e density, angle = 90, drawn correctly

trigonal bipyramidal

5 areas of e density, 120 and 90 angles, drawn correctly