View
231
Download
2
Category
Tags:
Preview:
Citation preview
Chapter 6.2 and 6.5Chapter 6.2 and 6.5Chapter 6.2 and 6.5Chapter 6.2 and 6.5
Covalent CompoundsCovalent Compounds
Covalent Bonds• Sharing Electrons
– Covalent bonds form when atoms share one or more pairs of electrons• nucleus of each atom is attracted to
electron cloud of other atom• neither atom removes an electron from
the other
Covalent Bonds• Sharing Electrons
– Covalent bonds• space where electrons move is called
molecular orbital– made when atomic orbitals overlap
Covalent Bonds
• Energy and Stability– Noble gases are stable (full octet)
(low P.E.)– Other elements are not stable (high
P.E.)• covalent bonding decreases potential
energy because each atom achieves electron configuration like noble gas
Covalent Bonds
• Energy and Stability•because P.E. decreases when atoms bond, energy is released –i.e., atoms lose P.E. when they bond
–loss of P.E. implies higher stability
Covalent Bonds• Energy and Stability
•potential energy determines bond length
–at minimum P.E., distance between two bonded atoms is called bond length
»bonded atoms vibrate»therefore, bond length is an
average length
Covalent Bonds• Energy and Stability
•bonds vary in strength–bond energy is the amount of
energy required to break the bonds in 1 mol of a chemical compound
–bond energy predicts reactivity–bond energy is equal to loss of
P.E. during formation
Bonds and Energy• Single bonds are the longest bonds
with the least bond energy• Double bonds are shorter, stronger
and have intermediate bond energy• Triple bonds are the shortest,
strongest, and have the highest bond energy
Covalent Bonds• Electronegativity
– Atoms share electrons equally or unequally•nonpolar covalent bond:
bonding electrons shared equally•polar covalent bond: shared
electrons more likely to be found around more electronegative atom
Covalent Bonds• Electronegativity
– Atoms share electrons equally or unequally•difference in electronegativity can be used to predict type of bond (but boundaries are arbitrary)
Practice: Calculate the bond type
• N and H• F and F• Ca and Cl• C and O
• Polar • Non-polar• Ionic• Polar
Covalent Bonds• Electronegativity
– Polar molecules have positive and negative ends• such molecules called dipoles• δ means partial in math and science• positive end—δ+
• negative end—δ-
• example: Hδ+Fδ-
Covalent Bonds• Electronegativity
– Polarity is related to bond strength•greater electronegativity
means •greater polarity
means •greater bond strength
Covalent Bonds• Electronegativity
– Bond type determines properties of substances•metallic bonds: electrons can move
from one atom to another—good conductors
•ionic bonds: hard and difficult to break apart
•covalent bonds: low melting/boiling points
Polarity of Molecules
• Two atoms: bond polarity is the molecular polarity
• More than 2 atoms: the geometry of the molecule must be considered
• If the bonds are non-polar, the molecular is non-polar
• Some molecules with polar bonds • can be non-polar
More • Sometimes the partial charges
cancel each other out because they are directly opposite each other
• Consider CO2 and CCl4• The symmetrical distribution of the
bonds leads to cancellation of the charges
Drawing and Naming• Lewis Electron-Dot Structures
– Lewis structures represent valence electrons with dots•position of electrons is symbolic
(not literal)•shows only the valence electrons
of an atom•dots around atomic symbol
represent electrons
Drawing and Naming• Lewis Electron-Dot Structures
– Drawing
–1. Gather information»draw Lewis structure for each
atom in compound; place one electron on each side before pairing
»determine total number of valence electrons
Drawing and Naming• Lewis Electron-Dot Structures
– Drawing•2. Arrange atoms
–arrange structure to show bonding
–halogens and hydrogen usually make one bond at end of molecule
–carbon usually in center
Drawing and Naming• Lewis Electron-Dot Structures
– Drawing•3. Distribute the dots so that each
atom satisfies octet rule (except H, Be, B)
•4. Draw the bonds as long dashes•5. Verify the structure by counting
number of valence electrons
Drawing and Naming• Lewis Electron-Dot Structures
– Polyatomic Ions• use brackets [] to show overall charge• example:
Drawing and Naming• Lewis Electron-Dot Structures
– Multiple Bonds• sharing two pairs of electrons is a
double bond• sharing three pairs of electrons makes
triple bonds• example:
Drawing and Naming• Lewis Electron-Dot Structures
– Resonance Structures•sometimes, multiple structures are possible
•show all possibilities•example:
Drawing and Naming• Naming Covalent Compounds
– First name: name of first element in formula• usually least electronegative• requires a prefix if more than one of them
– Second name: ends in –ide• requires a prefix if more than one of them
Practice Naming1. antimony tribromide
2. hexaboron silicide
3. chlorine dioxide
4. iodine pentafluoride
5. dinitrogen trioxide
6. ammonia
Molecular Shapes
• Determining Molecular Shapes– Three-dimensional shape helps
determine physical and chemical properties
– valence shell electron pair repulsion (VSEPR) theory predicts molecular shapes• based on idea that electrons repel one
another
Molecular Shapes
• Lewis structures show which atoms areconnected where, and by how many
bonds, but they don't properly show 3-D shapes of molecules.
To find the actual shape of a molecule, first draw the Lewis structure, and then use VSEPR Theory.
MOLECULAR MOLECULAR GEOMETRYGEOMETRY
VSEPRVSEPR • VValence alence SShell hell EElectron lectron PPair air
RRepulsion theory.epulsion theory.• Most important factor in determining Most important factor in determining
geometry is relative repulsion between geometry is relative repulsion between electron pairs.electron pairs.
MOLECULAR MOLECULAR GEOMETRYGEOMETRY
Molecule Molecule adopts the adopts the shape that shape that minimizes minimizes the electron the electron pair pair repulsions.repulsions.
VSEPR Rules• To apply VSEPR theory:• 1: Draw the Lewis structure of the
molecule and identify the central atom
• 2: Count the number of electron charge clouds (lone and bonding pairs) surrounding the central atom.
• 3: Predict molecular shape by assuming that clouds orient so they are as far away from one another as possible
Bond Angles
• Lone-pairs of electrons behave as if they are slightly bigger than bonded electron pairs and act to distort the geometry about the atomic center so that bond angles are slightly smaller than expected:
Bond Angles
• Methane, CH4, has a perfect tetrahedral bond angle of 109.5°, while the H-N-H bond angle of ammonia, NH3, is slightly less at 107°, trigonal pyramidal
•
Bond Angles• The oxygen of water
has two bonded electron pairs and two nonbonded "lone" electron pairs giving a total VSEPR coordination number of 4.
• But the geometry is defined by the relationship between the H-O-H atoms and water is said to be "bent" or "angular" shape of 105°.
105°
Molecular Shapes• Determining Molecular Shapes and
angles.– Let’s try some.
• CO
• CO2
• BF3
• CH4
• CBr4
• PCl3
Recommended