Chemistry Unit 3 Section 2: Covalent Bonds

• Explain how nonmetals share electrons to form bonds

• Distinguish between single, double, & triple bonds

• Draw electron dot structures for molecules

• Describe the difference between polar and nonpolar covalent bonds

​Molecular Compounds

• Covalent Bonds: bonds between nonmetal atoms that share some or all of their valence electrons

• Molecule: A neutral group of atoms held together by one or more covalent bonds

  • Can be two or more of one atom, or two or more different atoms

    • ​Diatomic molecule: A molecule containing only two of the same atoms

​Covalent vs Ionic

• Covalent Bonds Share Electrons

  • ​Covalent Bonds Cooperate and Share

  • Two non-metals

  • Low Melting/Boiling Points

  • Bonds are strong, but not as strong as ionic

• Ionic Bonds involve one atom losing electrons an another gaining electrons

  • Ionic Bonds- It's all about me, I will Take what I want​

  • A metal and a non-metal

  • High Melting/Boiling Points

  • Conduct Electricity

​Molecular Compounds

• There are many different models to show molecules

• One model show the different orbitals of the atoms combining into hybridized orbitals to make different shapes

​The Octet Rule in Molecules

• When nonmetals bond together, they don't gain or lose electrons, but share

• In general, Atoms still want to have 8 electrons in their valence shell to be stable

  • Exceptions:

    • Hydrogen only needs 2

    • Boron Only Needs 6

    • Sulfur and Phosphorus can have up to 12

• The number of Valence Electrons relates to the Group numbers

  • Group 1 Has 1

  • Group 2 has 2

  • Group 13-17 has (group number-10)​

​Electron Dot Structure

• ​Electron-Dot Notation: shows only the valence electrons of an atom, indicated by dots around the element symbol

  • Used to show the shared electrons in a molecule

  • Each Bond between atoms is made up of at least 2 electrons

    • More electrons in a bond make a bond stronger/harder to break

    • Different types of Bonds

      • Single Covalent Bond: a bond made of one pair of electrons

      • Double Covalent Bond: A bond made of two pairs of electrons

      • Triple Covalent Bond: A bond Made of three pairs of electrons

​Electron Dot Structure Rules

• ​If Carbon is present, it always goes in the middle

• Hydrogen and Halogens go on the outside

• Start placing your electrons and filling the outer atoms before the central atoms​

• ​One Dot shows a single electron

• A bar shows two electrons​

​Electron Dot Structure Sample 1

• ​Draw an electron dot structure for the diatomic molecule H₂

  • Step 1: Identify the number of Valence electrons in all atoms of the molecule

    • Hydrogen is in Group 1 and has 1 valence electron

    • We have 2 Hydrogens, so we have 2 valence Electrons Total

  • Step 2: Write the Elemental Symbols with a bit of space between them

  • Step 3: Start placing your electrons, starting with a bond formed by TWO Electrons

  • Step 4: Once all electrons are placed, Check to see that all atoms have 8 electrons around them

    • Two if it is hydrogen

​Electron Dot Structure Sample 2

• ​Draw an electron dot structure for the diatomic molecule H₂O

  • Step 1: Identify the number of Valence electrons in all atoms of the molecule

​Electron Dot Structure Sample 2

• ​Draw an electron dot structure for the diatomic molecule H₂O

  • Step 1: Identify the number of Valence electrons in all atoms of the molecule

    • Hydrogen is in Group 1 and has 1 valence electron

    • Oxygen is in Group 16 and has 6 Valence Electrons

    • Total amount of Valence Electrons: 1+1+6=8

  • Step 2: Write your elemental symbols

    • Hydrogen can never go in the middle!

  • Step 3: Connect the atoms with single atomic bonds, then place other electrons so that each atom has a full number of electrons

  • Step 4: Double Check Your Work

​Electron Dot Structure Sample 3

• ​Draw an electron dot structure for the diatomic molecule HOCl

  • Step 1: Identify the number of Valence electrons in all atoms of the molecule

    • Hydrogen is in Group 1 and has 1 valence electron

    • Oxygen is in Group 16 and has 6 Valence Electrons

    ​

​Electron Dot Structure Sample 3

• ​Draw an electron dot structure for the diatomic molecule HOCl

  • Step 1: Identify the number of Valence electrons in all atoms of the molecule

    • Hydrogen is in Group 1 and has 1 valence electron

    • Oxygen is in Group 16 and has 6 Valence Electrons

    • Chlorine is in Group 17 and has 7 Valence Electrons

    • Total amount of Valence Electrons: 1+6+7=14

  • Step 2: Write your elemental symbols

  • Step 3: Connect the atoms with single atomic bonds, then place other electrons so that each atom has a full number of electrons

  • Step 4: Double Check Your Work

​Electronegativity & Bonding

• ​In covalent bonds, both atoms pull on electrons

  • The Pull is not equal between the two electrons

    • Think Tug-of-War​

• Electronegativity: The ability of an atom to attract electrons towards itself in a chemical bond

  • A Higher electronegativity means an atom can pull electrons to it more

• Electronegativity follows a trend where it increases as you move across the period, and decreases as you move down

  • Fluorine is the most electronegative element with a value of 4.0​

​Electronegativity & Bond Types

• Covalent bonds fall into two categories depending on how much each atom pulls

  • Non-polar Covalent Bonds:​ a bond where electrons are shared "equally"

    • Electronegativity is less than 0.4

  • Polar Covalent Bonds: Electrons are shared unequally with one side positive and the other negative

    • Electronegativity difference is between 0.4 and 2.0

  • If the electronegativity is greater than 2.0, the bond is ionic

​Electronegativity & Types of Bonds

​Calculating Electronegativity Differences Sample Problem 1

• Calculate the Electronegativity difference in the bonds between K and F, and determine if it is polar or non-polar

  • Step 1: Find the electronegativity of​ the two elements

    • K=0.8

    • F=4.0

  • Step 2: subtract the smaller value from the larger value

    • 4.0-0.8

​Calculating Electronegativity Differences Sample Problem 1

• Calculate the Electronegativity difference in the bonds between K and F, and determine the bond type.

  • Step 1: Find the electronegativity of​ the two elements

    • K=0.8

    • F=4.0

  • Step 2: subtract the smaller value from the larger value

    • 4.0-0.8=3.2

  • Step 3: Compare to the values for polar and non-polar to determine the bond type

​Calculating Electronegativity Differences Sample Problem 2

• ​Use the electronegativity differences to classify bonding between sulfur (S) and Hydrogen

  • Step 1: Find the electronegativity of​ the two elements

    • S=2.5

    • H=2.1

  • Step 2: subtract the smaller value from the larger value

​Calculating Electronegativity Differences Sample Problem 2

• ​Use the electronegativity differences to classify bonding between sulfur (S) and Hydrogen

  • Step 1: Find the electronegativity of​ the two elements

    • S=2.5

    • H=2.1

  • Step 2: subtract the smaller value from the larger value

    • 2.5-2.1=0.4