Compounds UNIT 4 CH 7, 8, 9

Introduction / review  An atom has protons and neutrons in its nucleus.  Electrons move in energy levels around the nucleus.  While the protons and neutrons stay in the nucleus, electrons can move freely from one atom to another.

Valence electrons  Valence electrons are the electrons that are in the highest energy level of an atom.  These electrons are involved in forming bonds with other atoms.  Elements (except helium) have the same # of valence electrons as their group #.  Electron dot structures are used to show valence electrons.

 An electron dot formula (Lewis dot structure) of an element shows the symbol of the element surrounded by its valence electrons.  We use one dot for each valence electron.  Consider phosphorus, P, which has 5 valence electrons.

Practice  Give the valence electron dot structure for:  Mg  Cl CC  Ne

Octet rule  In order to become stable, atoms tend to either gain or lose valence electrons so that its highest energy level will become full with 8 electrons, similar to a noble gas. (except He, which has 2).  This is called the octet rule.

Ions  An atom is electrically neutral because it has the same # of protons (+) and electrons (-)  An atom becomes charged when it either gains or loses electrons.  A charged atom is called an ion

Ionic compounds  A cation (+) is formed when an atom loses electrons. Usually metals are cations.  An anion (-) is formed when an atom gains electrons. Usually nonmetals are anions  Cations and anions have opposite charges and are attracted to one another.  These attractive forces can hold the ions together in an ionic bond, forming a compound.

Ionic compounds  Ionic compounds are usually made up of a metal and a nonmetal.  Ionic compounds are often called: salts

Example of an ionic compound: NaCl

Properties of ionic compounds  High melting point  Low malleability – break and shatter easily  Can conduct electricity under certain conditions: molten or dissolved in water.

Formulas of ionic compounds  A chemical formula represents which & how many atoms are in a compound.  Cations and anions bond so that the compound has no net charge overall. The positive and negative charges cancel each other out. The positive ion is listed first in the formula.

Naming Ionic Compounds 1. Name the cation 1 st & the anion 2 nd. 2. Monatomic (1 element) cations use the element name without any change. 3. Monatomic anions use the root of their element name plus the suffix –ide, so chlorine become chloride, oxygen becomes oxide, etc.

Example: Na + Cl - NaCl: sodium chloride

Magnesium oxide: MgO Sodium oxide: Na 2 O

Trends in charge of ions  Group 1: always 1+  Group 2: always 2+  Al: always 3+  Group 7 (halogens): always 1-  Group 6: always 2-  Group 5 (N, P): always 3-

Practice  KBr  Since this has a metal (K) and a nonmetal (Br), we say it is an ionic compound.  So we name the positive ion – potassium and the negative ion with the ending changed to – ide, bromide.  Potassium bromide

Practice  Calcium Chloride  Again a metal and a nonmetal so it is ionic.  Calcium would form an ion with a 2+ charge  and chloride would be 1-.  Ca 2+ Cl -  in order for the compound to be neutral, how many Cl - would there need to be for every Ca 2+ ??  2 Cl - for every 1 Ca 2+  So the formula would be CaCl 2  The 2 next to Cl is an example of a subscript

Practice  Give the name of the following compounds: Na 2 O BaI 2 Give the formulas for: Magnesium bromide Aluminum fluoride Calcium nitride

Metal ions with more than one common charge  Certain metals can form more than one type of cation  Roman numeral (in between parenthesis) tells the charge on the cation Examples:  Iron (II) chloride, FeCl 2  Iron (III) chloride, FeCl 3

Some metal ions with more than one common charge  Fe: 2+ and 3+  Cu: 1+ and 2+  Sn: 2+ and 4+  Pb: 2+ and 4+  Co: 2+ and 3+  Mn: 2+, 6+ and 7+

Naming ionic compounds Certain ions are composed of several different atoms: polyatomic ions. If the compound has a polyatomic ion, simply name that ion without any change.

Polyatomic ions  NH 4 + Ammonium  OH - Hydroxide  NO 3 - Nitrate  CO 3 2- Carbonate  SO 4 2- Sulfate  PO 4 3- Phosphate  HCO 3 - Hydrogen carbonate / bicarbonate  C 2 H 3 O 2 - Acetate  SO 3 2- Sulfite

Practice  Na 2 CO 3  In this compound there are two ways to identify is as ionic.  First, it has a metal and a nonmetal.  Second, it has a polyatomic ion.  So we name the ions, positive ion first.  Sodium carbonate

More practice  FeCl 3 :  Cl is always 1- and there are 3 of those.  The compound is neutral, so Fe should have a charge of 3+  Iron(III) chloride

More practice  Copper(II) sulfate  Cu can be 1+ or 2+; here the Roman numeral tells you it is 2+  SO 4 is always 2-  CuSO 4

Practice  Magnesium Phosphate  magnesium – Mg 2+ ; phosphate – PO 4 3-  In order for the compound to be neutral we have to find the least common multiple between our two charges, 2 and 3. The LCM is 6.  2 goes into 6 – 3 times so Mg 3 ; 3 goes into 6 – 2 times so (PO 4 ) 2.  3 x +2 = +6 AND 2 x -3 = -6  Our compound is neutral.  Mg 3 (PO 4 ) 2  Use parenthesis around polyatomic ions when there are more than 1

Give the name for:  BaS  Barium sulfide  (NH 4 ) 2 SO 4  Ammonium sulfate  CoCO 3  Cobalt(II) carbonate

Give the formula for  Iron(III) nitrate  Fe(NO 3 ) 3  Tin(II) oxide  SnO  Lithium phosphate  Li 3 PO 4

Bonding in metals  Valence electrons of metal atoms can be modeled as a sea of electrons – they are mobile and can drift from one part of the metal to the other  Metallic Bond – the attraction of these “free-floating” electrons for the metal ions Metallic Bond  These bonds hold metals together and explain many of their physical properties

Molecular (covalent bonding) compounds  A covalent bond results from the sharing of electrons. The octet rule still applies  Covalent bonds generally occur when elements are close to each other on the periodic table.  The majority of covalent bonds form between nonmetallic elements. (remember that ionic bonds form between metals and non metals)

Example of a covalent bond

Octet Rule w/ covalent bonding

Diatomic Elements (elements that exist in pairs) *add this list to your chem friend  HydrogenH 2  OxygenO 2  NitrogenN 2  FluorineF 2  ChlorineCl 2  BromineBr 2  IodineI 2

Properties of molecular compounds  Usually have a lower MP (than ionic, metallic)  Generally soft  Non-conductors (in any state & aqueous)  Molecular Compounds can exist in all 3 states:  Solids – sugar, ice, aspirin  Liquids – water, alcohols  Gases – O 2, CO 2, and N 2 O (laughing gas)  Very important to organic chemistry, biochemistry, and pharmaceutics.

Naming Molecular (covalent bonding) Compounds 1. The 1 st element is named first, using the entire element name 2. The 2 nd element is named using the root of the element & adding the suffix –ide 3. Prefixes are used to indicate the # of atoms of each type that are present. *Exception: The 1 st element in a formula never uses the prefix mono-

Prefixes in Molecular (covalent) Compounds # of atoms Prefix# of atoms Prefix 1mono-6hexa- 2di-7hepta- 3tri-8octa- 4tetra-9nona- 5penta-10deca-

Examples NH 3 nitrogen trihydride (ammonia) N 2 H 4 dinitrogen tetrahydride H 2 Odihydrogen monoxide (common name water)

Practice  CO 2  Carbon dioxide  CO  Carbon monoxide  PCl 5  Phosphorus pentachloride N2O5N2O5  Dinitrogen pentoxide  NOTE – Be sure to drop the last vowel of the prefix if there would be any a-o, o-o, or a-a combinations (pentoxide, not pentaoxide).

Practice  Dichlorine monoxide  Cl 2 O  Tetraiodine nonoxide I4O9I4O9  Sulfur hexafluoride  SF 6

Ionic vs. Covalent  Look at the 1 st element:  1 st element metal: ionic compound  1 st element non-metal: covalent compound  Exception: ammonium (NH 4 + ): ionic compound

Practice  AlCl 3 : Al is a metal, so ionic name  Aluminum chloride  SiCl 4 : Si is a non-metal, so covalent name  Silicon tetrachloride