Chapter 7

 Ions are charged atoms.  An neutral atom becomes an ion when it either gains or loses electrons.  The imbalance between the number of protons and the number of electrons results in a charge on the atom.  An ion with a positive charge is called a cation.  An ion with a negative charge is called an anion.

 The octet rule states that atoms tend to lose, gain, or share electrons in order to acquire a full set of eight valence electrons.

 The electrostatic force that holds two ions together is called an ionic bonds.  compounds that contain ionic bonds are called ionic compounds.  Na + Cl  Na + + Cl -  [Ne] 3s 1 + [Ne] 3s 2 3p 5  [Ne] + [Ar]

 Most ionic compounds have a crystal like structure.  This structure is known as a crystal lattice.

 When ionic compounds are dissolved in water they break apart into a cation and an anion.  Not all ionic compounds dissolve in water.  Ones that do are called electrolytes.

 The chemical formula for an ionic compound, called a formula unit, represents the simplest ratio of ions involved.  MgCl 2  A monatomic ion is a one-atom ion.

 Representative elements of the same groups usually have the same ionic charges.  Group 1: H +, Li +, Na +  Group 2: Be 2+, Mg 2+, Ca 2+  Group 15: N 3-, P 3-, As 3-  Group 16: O 2-, S 2-, Se 2-  Group 17: F -, Cl -, Br -  Transition elements can usually have more than one charge. Example Fe 2+ and Fe 3+

 Most transition metals can have a few different charges.  Example:  Fe 2+, Fe 3+  Fe 2 O 3  FeO

 Polyatomic ions are ions that are made up of more than one atom.  There are a few you need to memorize….sorry.  Ammonium: NH 4 +  Nitrite: NO 2 -  Nitrate: NO 3 -  Hydroxide: OH -  Hypochlorite: ClO -

 Chlorite: ClO 2 -  Chlorate: ClO 3 -  Perchlorate: ClO 4 -  Carbonate: CO 3 2-  Sulfite: SO 3 2-  Sulfate: SO 4 2-  Peroxide: O 2 2-  PO 4 3-

 The cation comes first in the name then the anion. If the anion is a monatomic ion we give it the ending –ide.  NaCl – Sodium Chloride  When the formula unit contains two or more of the same polyatomic ion we put it in parentheses with a subscript to show how many ions are present.  Ba(NO 3 ) 2 – Barium Nitrate  If the cation is a transition element the charge of the cation is written as roman numerals in parentheses.  Co(SO 4 ) – Cobalt (II) Sulfate

 NH 4 Cl  Fe(NO 3 ) 3  TiBr 3  Pb(SO 4 ) 2  Chromium (VI) Phosphate  Tin (II) Nitrate  Cobalt (III) Oxide  Chromium (III) Hydroxide

 Barium Carbonate  Aluminum Hydroxide  Copper (I) Sulfide  Lead (II) Phosphate  Zinc (II) Iodide  Chromium (III) Sulfite  NH 4 Br  NaClO 3  Fe 2 S 3  AgNO 3  CuF 2  Ni(ClO) 3  Ru 3 (PO 4 ) 4

Chapter 8

 Why do atoms bond?  Atoms bond to become more stable.  Atoms are most stable when they have 8 valence electrons. (i.e. the same electron configuration as the closest noble gas.)  A chemical bond that results from two atoms sharing electrons is called a covalent bond.

1. The first element in the formula is always named first, using the entire name of the element. 2. The second element in the formula is named using its root and adding the ending –ide. 3. Prefixes are used to indicate the number of atoms of each element that are present in the compound.  NO 2 – Nitrogen Dioxide  N 2 O 2 – Dinitrogen Dioxide

 CO 2  SO 2  NF 3  CCl 4  What is the formula of Diarsneic trioxide?

 Fluorine, and some other common elements, usually exist as a molecule of F 2.  Fluorine’s electron configuration is….  Fluorine’s structure is…

1. Add up all the valence electrons involved. 2. Determine which atom is the central atom. 3. Form single bonds between the central atom and all of the surrounding atoms. (each single bond uses 2 valence electrons) 4. Place electrons around the surrounding atoms until all have eight or until we have run out of valence electrons. 5. If there are any remaining valence electrons place them on the central atom.

 PH 3  HCl  PBr 3  CCl 4  SiH 4  NH 3  NF 3

 When only one pair of electrons are shared (like in F 2 ) it is called a single bond.  Single bonds are also called sigma (σ) bonds  Group 17: ◦ The halogens all have seven valence electrons and will therefore be able to form one covalent bond.  Group 16: ◦ The elements in group 16 have six valence electrons and can therefore make two single covalent bonds.  Group 15: ◦ The elements in group 15 have five valence electrons and can therefore make three single covalent bonds.  Group 14: ◦ The elements in group 14 have four valence electrons and can therefore make four single covalent bonds.

 Some elements are able to share more than one pair of electrons in order to obtain eight valence electrons.  A double covalent bond forms when two atoms share two pairs of electrons.  A triple covalent bond forms when two atoms share three pairs of electrons.  One bond in a multiple bond molecule is still called a sigma bond and the second and/or third are called pi (π) bonds.

 To draw a lewis structure of a polyatomic ion we first figure out how many valence electrons we have, then add or subtract electrons to account for the charge.  SO 4 2-

 The strength of a covalent bond depends on the distance between the bonded nuclei. F2F2 ◦ Single covalent bond ◦ Bond length 1.43 x m O2O2 ◦ Double covalent bond ◦ Bond length 1.21 x m N2N2 ◦ Triple covalent bond ◦ Bond length 1.10 x m

 A small group of molecules have have an odd number of valence electrons and be unable to form an octet around each atom.  NO 2  An other group of molecules have expanded octets.  SF 6

 Sometimes its possible to have more than one correct lewis structure for a molecule.  In these cases all of the possible lewis structures are called resonance structures.  NO 3 -  SO 2 O3O3

 Binary acids (meaning molecules with just hydrogen bonded to another single atom)  The first word has the prefix – hydro – the rest of the first word consists of a from of the name of the second atom followed by the prefix – ic.  HCl – hydrochloric acid  HF – hydrofluoric acid  HI – hydroiodic acid

 Oxyacids are acids that contain hydrogen bonded to a polyatomic anion.  The first word of the name of these acids consists of the root of the polyatomic ion. If the polyatomic ion’s name ends in –ate you replace it with the ending –ic. If the polyatomic’s name ends in –ite you replace it with –ous.  HClO 3 – Chloric acid  HClO 2 – Chlourous acid  HNO 3 – Nitric acid  HNO 2 – Nitrous acid

 HI  HClO 3  H 3 PO 4  H 2 SO 4  H 2 S

 The molecular geometry, or shape, of a molecule can be determined once a lewis structure is drawn.  The model used to determine a molecules shape is the VSEPR model  V = Valence  S = Shell  E = Electron  P = Pair  R = Repulsion

 Each bond, or set of nonbonding electrons, repels the other electrons in a molecule.  Molecules form the shape that results in the smallest interactions between the electrons in the molecule.  The shape of a molecule is dependant on what the eight, or sometimes more, electrons around the central atom are doing.

 A covalent bond forms between two atoms when a pair of electrons occupies the space between the atoms. ◦ This is a bonding pair of electrons ◦ The region these electrons occupy is called the electron domain.  A nonbonding pair of electrons (Lone pair) reside on only one atom.  Example NH 3 : How many bonding pairs? How many nonbonding pairs?

 VSEPR predicts that the best arrangement of electron domains is the one that minimizes the repulsion among them. ◦ The arrangement of electron domains about the central atom of an Ab n molecule is it’s electron domain geometry.  There are five different electron domain geometries. ◦ Linear, Trigonal planar, Tetrahedral, trigonal bipyramidal, and octahedral.

 Consider CH 4, NH 3 and H 2 O  All three have the same electron domain geometry but have different molecular geometry and bond angles.  Electron domains for nonbonding electron pairs exert greater repulsive forces on other electron domains. This is why the bond angles decrease as the number of nonbonding electron domains increases.

 The type of bond formed during a reaction is related to each atoms attraction for electrons.  Electron affinity is a measure of the tendency of an atom to accept and electron.  Electron affinity increases as you move from left to right on the periodic table and from bottom to top.  The difference between electronegativity determines what type of bond is formed.

Electronegativity DifferenceBond Type > 1.7Ionic Bond 0.4 – 1.7Polar Covalent Bond < 0.4Nonpolar Covalent Bond Shaken not stirredJames Bond

 Because not all atoms attract electrons equally sometimes we have one atom in a bond being “greedy”.  HCl  Electronegativity of Cl – 3.16  Electronegativity of H – 2.20  Difference – 0.96

1. Fe(OH) 2 2. PCl 5 3. Sulfurous Acid 4. BF 3 5. SO 2 6. Nitrogen Trihydride 7. Ni 2 (CO 3 ) 3 8. Hydrobromic Acid 9. NO Carbon Monoxide 11. HNO HClO 4