Chemical Bonding & Molecular Shapes

Electrons The number of valence electrons determines many of the chemical properties an element has Elements in the same group have similar chemical and physical properties because they have the same number of valence electrons

Electrons Group 1 = 1 valence electron Group 2 = 2 valence electrons *skip the middle* Group 13 = 3 valence electrons Group 14 = 4 valence electrons Group 15 = 5 valence electrons Group 16 = 6 valence electrons Group 17 = 7 valence electrons Group 18 = 8 valence electrons

Lewis Dot Structures Valence electrons can be represented as a series of dots surrounding an atomic symbol This is called an electron-dot structure, or more commonly, a Lewis dot structure Lewis dot structures immediately tell you two things about an element: How many valence electrons the element has How many of those electrons are paired

C B O Lewis Dot Structures Completing Lewis dot structures must be done in a specific way in order to correctly show paired and unpaired electrons After identifying the number of valence electrons an atom has: Write down the element symbol. Fill in the electrons ON EACH SIDE of the symbol until the correct number of electrons is reached For example – Carbon = 4 valence electrons Boron = 5 valence electrons Oxygen = 6 valence electrons B O

Paired & Unpaired Electrons Paired valence electrons are relatively stable Do not usually form chemical bonds Referred to as nonbonding pairs Unpaired valence electrons are much more likely to participate in chemical bonding They become paired with electrons from another atom

Ions Review: when an atom gains or loses electrons, there is a net electric charge If electrons are lost, there are more protons than electrons—the ion’s charge is positive Called a cation Metals are more likely to lose electrons If electrons are gained, there are more electrons than protons—the ion’s charge is negative Called an anion Nonmetals are more likely to gain electrons

Ions Atoms tend to lose or gain electrons in order to have a full outermost shell (usually 8 electrons) Noble gases already have 8 electrons in their outermost orbital, they are inert Alkali metals (group 1) and Halogens (group 17) are very reactive because they are each only 1 electron away from having a full outermost shell

Ionic Bonding Ionic bonds results from a transfer of electrons Happens between a metal and nonmetal When a metal atom (wants to lose) is placed in contact with a nonmetal (wants to gain)—electrons are transferred and two oppositely charged ions are formed The two ions are attracted to each other by the electric force of attraction

Ionic Bonding Compounds with ionic bonds are called ionic compounds The properties of these compounds are completely different from the elements that form them

Naming Ionic Compounds Identify the names of the elements in the compound Identify which is a metal and which is a nonmetal Write the name of the metal with a capital letter Write the name of the nonmetal with a lowercase letter Change the ending of the nonmetal to –ide EXAMPLES Sodium and Chorine Lithium and Fluorine

Total positive charge = total negative charge Ionic Compounds Ions have different charges Ionic compounds want to have an overall charge of zero Total positive charge = total negative charge From the charges on the ions, you can determine the formula of an ionic compound

Determining Formula of Ions Example: BORON OXIDE Write the symbols of both elements (cation 1st, anion 2nd) Write the valence of each as a superscript Drop the positive and negative signs Crisscross the superscripts so they become subscripts Reduce when possible (not possible here) B O B3+ O2- B3 O2 B3 O2 B2O3

Transition Metals Transition metals have a varying number of valence electrons For example: Copper can form Cu+ and Cu2+ ions These are written as Copper (I) and Copper (II) Iron can form Fe2+ or Fe3+ ions These are called Iron (II) and Iron (III) The Roman numeral represents the CHARGE of the ion

Transition Metals You can use the ionic formula to determine the charge of the transition metal ion If you are given CuCl2, what would be the charge on the copper atom? Copper would have a 2+ charge because of the “2” subscript after chlorine

Polyatomic Ions Ions that have many atoms in one ion You cannot break these ions apart—they always appear together Ammonium NH4+ Chlorate ClO3- Peroxide O22- Acetate CH3CO2- Perchlorate ClO4- Chromate CrO42- Nitrate NO3- Permanganate MnO4- Dichromate Cr2O72- Nitrite NO2- Carbonate CO32- Silicate SiO32- Hydroxide OH- Sulfate SO42- Phosphate PO43- Hypochlorite ClO- Sulfite SO32- Arsenate AsO43- Chlorite ClO2- Thiosulfate S2O32- Arsenite AsO33- Cyanide CN- Thiocyanate SCN- Borate BO33- Bicarbonate HCO3- Bisulfate HSO4- Bisulfite HSO3-

Polyatomic Ions DO NOT MEMORIZE THE LIST! Remember that polyatomic ions cannot be broken apart Example: If you have a sodium ion (Na+) and a hydroxide ion (OH-) they will combine to form NaOH (read: sodium hydroxide) If you see a name that ends in –ate, this is an indicator that a polyatomic ion is present

Covalent Bonds Atoms joined by covalent bonds share electrons Covalent bonds are usually formed between two nonmetals

Naming Covalent Compounds For covalent compounds of two elements, numerical prefixes are used to tell how many atoms of each element are in the molecule If there is only one atom of the first element, the name does not get a prefix The element that is farther to the right in the periodic table is named second, and the ending is changed to -ide # of Atoms Prefix 1 mono- 2 di- 3 tri- 4 tetra- 5 penta- 6 hexa- 7 hepta- 8 octa- 9 nona- 10 deca-

Covalent Compounds In a Lewis dot structure, a straight line represents electrons being shared between the atoms A single straight line = two electrons (each atom contributing one) Electrons being shared are called the bonding pairs Electrons that are not a part of the bond are called nonbonding pairs

Covalent Compounds Covalent compounds can also have more than one bond between them Double bonds occur when atoms share 4 electrons— shown by two straight lines Triple bonds occur when atoms share 6 electrons— shown by three straight lines

Lewis Dot Structures Sum the valence electrons from all atoms Every atom wants to have 8 electrons around it 8 electrons means it has a full octet, or outermost orbital, which makes the atom stable Exception is hydrogen—will only have two electrons To draw a dot structure for a compound – Sum the valence electrons from all atoms Use a pair of electrons to form a bond between each pair of bound atoms Arrange the remaining electrons to satisfy the octet rule (or duet rule)

Write the Lewis dot structure for: H2O C2H6 H2O2

VSEPR Theory Stands for valence-shell electron pair repulsion In 3-D, all atoms and electrons will try to stay as far away from each other as they possibly can Electrons and the electron shells of other atoms are all negative, so similar charges repel one another Remember: electrons that are shared between atoms are called bonding electrons; nonbonding electrons are also called lone pairs

Molecular Shape VSEPR theory can predict the geometry of individual molecules using the number of electron pairs surrounding the central atom

PhET Colorado Simulation Molecular Shape 1. Draw the Lewis dot structure 2. Determine the number of bonded pairs on the CENTRAL atom 3. Determine # of lone pairs on the CENTRAL atom 4. ADD these (STERIC #) 5. Use steric #, # of lone pairs, and the table to determine shape PhET Colorado Simulation

Lone Pairs Steric #2 Steric #3 Steric #4 Steric #5 Steric #6 Linear Linear Trigonal Planar Tetrahedral Trigonal Bipyramidal Octahedral 1 (Linear) (Bent) Trigonal Pyramidal Seesaw Square Pyramidal 2 n/a Bent T-Shaped Square Planar 3

Draw the Lewis dot structure and determine its molecular geometry: CCl4

Draw the Lewis dot structure and determine its molecular geometry: NI3

Draw the Lewis dot structure and determine its molecular geometry: SO2

Draw the Lewis dot structure and determine its molecular geometry: H2O