Chapter 4 Flashcards
(24 cards)
Ionic bonding
Metals combine with non-metal
Electrons in the outer shell of the metal atoms are transferred to the non-metal atoms
Non-metal atoms gain electrons to fill their outer shell
Metal and non-metal end up with an electronic configuration of a noble gas
Dot-and-cross diagram
Draw magnesium oxide and calcium chloride
Outer electron shells only
The charge of an ion is spread evenly using square brackets
Charge on each ion written at the top right-hand corner of the square brackets
Covalent bonding
2 non-metal atoms combine they share one or more pairs of electrons
Lone pairs
Pairs of outer-shell electrons not used in bonding
Dative covalent bonding
Coordinate bonding
Formed when one stone provides both the electrons needed for a covalent bond
Requirements:
One atom have a lone pair of electrons
A second atom having an infilled orbital to accept the lone lair, an electron-deficient compound.
Bond length and bond strength
Bond length: multiple bonds are short because they have a greater quantity of negative charge between the two atomic nuclei
Bond energy: energy needed to break one mole of a given bond in a gaseous molecule influences reactivity because the higher the bond energy the more energy is needed to overcome one mole of a given bond
Electron pair repulsion theory
All electrons are all negatively charged so they repel each other when they are close to together, pair of electrons in the bonds surrounding the central atom will repel other electrons pairs this repulsion forces the pairs of electrons apart until repulsive forces are minimized
What do the shape and bond angles of a covalently bonded molecule depend on
Number of pairs of electrons around each atom
Whether these pairs are lone pairs or bonding pairs
Compare the different types of electron pairs
Lone pairs: electron charge clouds are more concentrated, wider and slightly closer to the nucleus of the central atom than bonding pairs
Order of repulsion
Lone pair - lone pair > lone pair - bond pair> bond pair-bond pair
What is the shape and angle of
2 lone pairs
2 bonding pairs
3 elements
Shape: angular
Angle: 104.5
What is the shape and angle of
0 lone pairs
4 bonding pairs
4 elements
Shape: tetrahedral
Angle: 109.5
What is the shape and angle of
0 lone pairs
2 bonding pairs
3 elements
Shape: linear
Angle: 180
What is the shape and angle of
0 lone pairs
3 bonding pairs
3 elements
Shape: trigonal planar
Angle: 120
What is the shape and angle of
1 lone pairs
3 bonding pairs
3 elements
Shape: pyramidal
Angle: 107
What is the shape and angle of
0 lone pairs
5 bonding pairs
5 elements
Shape: trigonal bipyramid
Angle: 120 within the plane, 90 above the plane
What is the shape and angle of
0 lone pairs
6 bonding pairs
7elements
Shape: octahedral
Angle: 90
Molecular orbital
Single covalent bond is formed when 2 non-metal atoms combine
Each atom that combines has an atomic orbital containing a single impaired electron
Atomic orbitals overall so that a combined orbital is formed containing two electrons when forming a covalent bond
Amount of overlap of atomic orbitals determines the strength of the bond
Sigma bonds
Symmetrical electron density about a line joining the nuclei of the atoms forming the bond
When hybridized orbitals overlap linearly
Hybridization
P atomic orbitals overlap linearly to form covalent bonds
When involved in forming single bonds they include some s orbital character
The orbital is slightly altered in shape to make one of the lines of the p orbital buffer
In sp3 each orbital has a third s character and two thirds p character
Pi bonds
Bonds formed by the sideways overlap of p orbitals
Asymmetrical about the axes joining the nuclei of the atoms forming the bond
Shape of ethane and ethene
Ethane:
All sigma bonds, equal repulsion, angle throughout 109.5
Ethene:
3 of its 4 outer electrons to form sigma bonds
2 are formed with hydrogen atoms
1 is formed with the other carbon atom
4th electron form each carbon atom occupies a p orbital which overlaps sideways forming a pi bond
Planar molecule for maximum overlap of the p orbitals that form pi bond
Electron clouds of pi bonds above and below the plane
120 in H-C-H because 3 areas of electron density of sigma bonds are equally distributed
117 because of pi bond to minimizes repulsive forces
Metallic bonding
Atoms are packed closely together in a regular arrangement called a lattice
Metal atoms in a lattice tend to lose their outer shell electrons and become cations
Outer shell electrons occupy new energy levels and are free to move throughout the metal lattice
These are delocalized electrons, they’re no associated with any one particular atom or bond
Metallic bonding is strong because ions are held together by strong electrostatic attraction between their positive charges and the negative charges of the delocalized electrons which acts in all directions
What does the strength of metallic bond increase by
Increasing positive charge on the ions in the metal lattice
Decreasing size of metal ions in the lattice
Increasing number of mobile electrons per atom
Metallic bonding and the properties of metals
High melting/boiling points: it takes a lot of energy to weaken the strong attractive forces between the metal ions and delocalized electrons
Mecury: liquid because some of the electrons in a mercury atom are bound more tightly than usual to the nucleus weakening the metallic bonds between atoms
Conductivity: current flows because the delocalized electrons are free to move and conduct heat because of vibrations passed on from one metal ion to the next
