Covalent Compounds

Two Types of Bonds Ionic: Electrons are transferred Covalent: Electrons are shared Non-polar covalent: equally shared Polar Covalent: unevenly shared

Bond Polarity

Review: What is electronegativity? ability of an atom to attract electrons Which element is the most electronegative? Fluorine - Has 7 valence e- and wants 8 H F

Polar bond : covalent bond with greater electron density around one of the two atoms H F electron poor region electron rich region e- poor e- rich F H d+ d-

1 18 2 13 14 15 16 17 3 4 5 6 7 8 9 10 11 12

Electronegativity Difference What type of Bond is it? Electronegativity Difference Bond Type 0 to 0.3 Nonpolar Covalent 0.4 to 1.6 Polar Covalent  1.7 Ionic Increasing difference in electronegativity Nonpolar Covalent share e- Polar Covalent partial transfer of e- Ionic transfer e-

Classify the following bonds as ionic, polar covalent,or covalent: Cs to Cl Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic H to S H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent Cl to N Cl – 3.0 N – 3.0 3.0 – 3.0 = 0 Nonpolar Covalent

Do you notice a pattern for the combo of elements that are ionic vs covalent? Ionic bonds form between: Covalent bonds form between: Identify the following as ionic, covalent, or both: CaCl2 BaSO4 CO2 AlPO4 SO3 H2O

Properties of Covalent Compounds Usually soft and squishy Not soluble in water Does not conduct electricity Soluble in organic solvents Low melting points Low boiling points

Properties of Ionic Compounds Combination of ions (cation/anion) Tightly packed solids in a crystal lattice Hard and Brittle Usually soluble in water Conducts electricity when dissolved High melting points High boiling points

Naming Covalent Compounds

NAMING COMPOUNDS Nonmetal – Nonmetal USE PREFIXES! Change the ending of the second word to -ide No mono on the first word Drop any double vowels

Covalent Prefixes Number of Atoms Prefix 1 Mono- 2 Di- 3 Tri- 4 Tetra- 5 Penta- 6 Hexa- 7 Hepta- 8 Octa- 9 Nona- 10 Deca-

Examples CO CO2 SO2 SO3 N2H4 N2O3 Carbon Monoxide Carbon Dioxide Sulfur Dioxide Sulfur Trioxide Dinitrogen Tetrahydride Dinitrogen Trioxide

Examples Si2F6 C3Cl8 disilicon hexafluoride tricarbon octachloride PBr5 NO SeF2 H2O disilicon hexafluoride tricarbon octachloride phosphorus pentabromide nitrogen monoxide selenium difluoride dihydrogen monoxide

EMPIRICAL AND MOLECULAR FORMULAS

Define Empirical Formula: A chemical formula that gives the simplest whole-number ratio of the elements in the formula. Which of the following is an empirical formula? CO2 C2O4 N2H4 NH2

Define Molecular Formula: A chemical formula that gives the actual number of the elements in the molecular compound. For the following molecular formulas, write the empirical formula: Molecular: Empirical: C2H4 C6H12O6 C9H21O6N3

Lewis Structures

This is called the octet rule Eight electrons in the valence shell (filling s and p orbitals) make an atom STABLE This is called the octet rule Bond formation follows the octet rule… Chemical compounds tend to form so that each atom: by gaining, losing, or sharing electrons, has an octet of electrons in its valence energy level.

Lewis Dot Diagrams an electron-configuration notation with only the valence electrons of an element are shown, indicated by dots placed around the element’s symbol. tracks the number of valence electrons the inner core electrons are not shown

Lewis Dot Practice Li N F Be O Ne

Lewis Structures for Compounds The pair of dots between two symbols represents a shared pair. How many shared pairs does each fluorine have below? An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.

Each dash represents TWO electrons Lewis Structures The shared pair of electrons is often replaced by a long dash. Each dash represents TWO electrons

Why should two atoms share electrons? To get a valence of 8 electrons! 7e- 7e- 8e- 8e- F F + F Lewis structure of F2 lone pairs F single covalent bond single covalent bond F

Multiple Covalent Bonds double bond: covalent bond in which two pairs of electrons are shared between two atoms shown by two side-by-side pairs of dots or by two parallel dashes

Multiple Covalent Bonds triple bond: covalent bond in which three pairs of electrons are shared between two atoms shown by three side-by-side pairs of dots or by three parallel dashes

Lengths of Covalent Bonds Bond Type Bond Length (pm) C-C 154 CC 133 CC 120 C-N 143 CN 138 CN 116 Bond Lengths Triple bond < Double Bond < Single Bond

Bond Length and Bond Energy As atomic size increases, bond length increases, and as a result bond energy decreases As you increase the number of bonds between two atoms, energy increases, while bond length decreases. Rubber band example

Bond Length and Bond Energy Examples Which bond is greater in length: Br2 or F2? The HF bond is 570 pm, the H2 bond is 436 pm, which bond requires more energy to break? Which bond would require more energy to break C-C single bond or C=C double bond? Which bond is longer?

Writing Lewis Structures Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. Complete an octet for all atoms except hydrogen If structure contains too many electrons, form double and triple bonds on central atom as needed.

Write the Lewis structure of nitrogen trifluoride (NF3). Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons F N

Write the Lewis structure of the carbonate ion (CO32-). Step 1 – C is less electronegative than O, put C in center Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e- 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e- 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24 O C

resonance structure: When there are two or more Lewis structures for a single molecule What are the resonance structures of the carbonate (CO32-) ion? O C - O C - O C -

Some elements do not follow the octet rule Be F B There can also be expanded octets!

Molecular Geometry

VSEPR THEORY Lewis Dot Diagrams are 2D but we live in a 3D world. How are molecules actually arranged?? Follows the Valance Shell Electron Pair Repulsion Theory or VSEPR

Number of Surround Atoms AB2 – Linear Number of Surround Atoms Number of Lone Pairs Bond Angle 2 180˚ Cl Be

Number of Surround Atoms AB3 – Trigonal Planar Number of Surround Atoms Number of Lone Pairs Bond Angle 3 120˚

Number of Surround Atoms AB2E1 – Bent Number of Surround Atoms Number of Lone Pairs Bond Angle 2 1 <120˚

Number of Surround Atoms AB4 – Tetrahedral Number of Surround Atoms Number of Lone Pairs Bond Angle 4 109.5˚

AB3E1 – Trigonal Pyramidal Number of Surround Atoms Number of Lone Pairs Bond Angle 3 1 107˚

Number of Surround Atoms AB2E2 – Bent Number of Surround Atoms Number of Lone Pairs Bond Angle 2 104.5˚

Predicting Molecular Geometry Draw Lewis structure for molecule. Count number of lone pairs on the central atom and number of atoms bonded to the central atom. Use VSEPR to predict the geometry of the molecule. What are the molecular geometries of SO2 and SF4? C F S O AB4 AB2E tetrahedral bent

Intermolecular Forces

Intermolecular forces: attractive forces between molecules. Intramolecular forces: attractive forces within a molecule (the bonds) Intermolecular Forces Intramolecular Forces Intramolecular Forces intramolecular forces are much stronger than intermolecular forces

Dipoles What is a dipole? A polar molecule Uneven sharing of electrons so there is a separation of charge

Dipole-Dipole Forces Attraction between two polar molecules — — + +

Hydrogen Bonding Special type of Dipole – Dipole Attraction between: Hydrogen and Nitrogen/Oxygen/Fluorine

Dipole – Induced Dipole Attraction between one polar and one nonpolar molecule — + Electrons shift toward positive end of dipole — — + +

London Dispersion Forces Attraction between two nonpolar molecules Electrons become uneven and form a dipole — — + +

Strength of IMF Hydrogen Bond Dipole – Dipole Dipole – Induced Dipole London Dispersion Forces strongest weakest

Which of the following molecules is polar? H2O, CO2, SO2, and CH4 O H S O dipole moment polar molecule dipole moment polar molecule C H C O no dipole moment nonpolar molecule no dipole moment nonpolar molecule

CH4 is nonpolar: dispersion forces. What type(s) of intermolecular forces exist between each of the following molecules? HBr HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules. CH4 CH4 is nonpolar: dispersion forces. S O SO2 SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

What does IMF effect? Viscosity Surface Tension Cohesion/Adhesion Boiling Point

Stronger IMF  Higher Viscosity Measures a fluid’s resistance to flow Stronger IMF  Higher Viscosity

Stronger IMF  Higher Surface Tension result of an imbalance of forces at the surface of a liquid. Stronger IMF  Higher Surface Tension

Stronger IMF  Higher Boiling Point Point at which liquid particles escape the surface of the liquid into the gas phase Stronger IMF  Higher Boiling Point

Adhesion and Cohesion Cohesion: intermolecular attraction between like molecules Adhesion: intermolecular attraction between unlike molecules Adhesion Cohesion