Unit II: Atoms Molecules & Ions A. B. C. D. The Periodic Table Why and how compounds form Rules for writing ions Writing formulas and naming compounds 1. 2. 3. Binary compounds Non-binary compounds Formula writing & naming of acids and bases E. Chemical Bonding
Two general groups of elements Metals Good conductors of heat & electricity Shiny, ductile, malleable Lose electrons to form positive ions Nonmetals Poor conductors Gain electrons to form negative ions Dull and without luster 2-A
Antimony Tin Arsenic Germanium Sulfur Iodine 2-A
Why and How Compounds Form e P N N P e 2-B
How Ionic Compounds Form An atom losing electrons become positively charged (cation) An atom gaining electrons becomes negatively charged (anion) Anions and cations attract each other and form compounds, lowering their total energy 2-B
Ionic Bonding Energy Energy e e e PNNP N PN e Li e e 1+ e NNN PPN PNNPPP N N P P e e Lithium e e e e F 1- Flourine 2-B
Ionic Bonding e e e e PNNPN N P e Li e NNN PPN PNNPPP N N P P e 1+ e e e e e F 1- Lower total energy 2-B
How Covalent Compounds Form e P P e Lower total energy 2-B
Metals Metals Cations Cations Ionic Ionic Compounds Compounds Nonmetals Nonmetals Anions Anions Nonmetals Nonmetals + Nonmetals Nonmetals Covalent Covalent Compounds Compounds Metals Metals + Metals Metals Homogeneous Homogeneous Mixtures Mixtures 2-B
Rules for Writing Ions Group IA Group IIA Group IIIA Alkali metals Alkaline earth metals Aluminum group +1 +2 +3 Group VIIA Group VIA Group VA Halogens Oxygen group Nitrogen group -1-2 -3 2-C
Writing Formulas and Naming Binary Compounds 1. Fixed oxidation state metals + nonmetals 2. Variable oxidation state metals + nonmetals 3. Nonmetals combined with nonmetals in covalent compounds 2-D
Binary ionic compounds (metals with fixed oxidation states) First name the metal then Use the root of the nonmetal with an ide ending Example. + - Na + Cl = NaCl Sodium + Chlorine = Sodium Chloride 2-D
Lecture Problems: II-2 (pg.61) Write the names of the following. magnesium fluoride A. MgF2 B. AlP aluminum phosphide C. Na2O sodium oxide 2-D
Lecture Problems: II-3 (pg.61-62) Use ionic charges to write the formulas of the following. A. lithium sulfide Li2S B. calcium nitride Ca3N2 C. barium chloride BaCl2 2-D
Binary ionic compounds (metals with variable states) First name the metal followed by its oxidation state in Roman Numerals inside parentheses Then add the root of the nonmetal with an ide ending Note: You must know the oxidation state of the metal to correctly name the compound! 2-D
Lecture Problems: II-4 (pg. 63) Given the following names, write the formulas A. iron (II) fluoride FeF2 B. lead (II) chloride PbCl2 C. chromium (III) sulfide Cr2S3 2-D
Lecture Problems: II-5 (pg. 63) Write the names of the following compounds: A. CuBr2 copper (II) bromide B. NiO nickel (II) oxide C. TiF4 titanium (IV) fluoride 2-D
Binary covalent compounds Name the first nonmetal using a prefix to indicate how many atoms are present Then name the second nonmetal using a prefix to identify quantity, and the ide ending Notes: 1. Drop the mono prefix from the element named first 2. Name the element located lower or to the left on the periodic table first 2-D
Covalent Compound Prefixes 1 = mono 2 = di 3 = tri 4 = tetra 5 = penta 6 = hexa 7 = hepta 8 = octa 9 = nona 10 = deca Drop the mono prefix from the first element named
Lecture Problem: II-6 (pg. 65) A. N2O4 dinitrogen tetraoxide B. P4S10 tetraphosphorus decasulfide Lecture Problem: II-7 (pg. 65) C. carbon dioxide CO2 D. trisulfur heptafluroide S3F7 E. phosphorus pentachloride PCl5 2-D
Non-Binary Compounds Generally consist of a cation with a polyatomic anion May also contain polyatomic cations with mono or polyatomic anions 2-D
Polyatomic ions with -1 charge OHCNNO3NO2C2H3O2MnO4HCO3HSO4- hydroxide cyanide nitrate nitrite acetate permanganate bicarbonate bisulfate ClOClO2- hypochlorite chlorite ClO3- chlorate ClO4- perchlorate See page 67 You will need to know them 2-D
Other Negative Polyatomic Ions CO3 2- carbonate SO4 2- sulfate SO3 2- sulfite CrO4 2- chromate Cr2O7 2- dichromate PO4 3- phosphate 2-D
Positive Polyatomic Ions NH4 + ammonium 2-D
Types of Compounds with Polyatomic Ions Metal of fixed oxidation state with polyatomic anion Metal with variable oxidation state with polyatomic anion Polyatomic cation with monoatomic anion Polyatomic cation with polyatomic anion 2-D
Metal of fixed oxidation state with a polyatomic ion Name the metal, then name the polyatomic ion You must make sure to balance the charges between ions If multiple polyatomic ions are needed, use () to describe how many are needed. 2-D
Examples Al2(SO4)3 aluminum sulfate Ca(OH)2 calcium hydroxide barium phosphate Ba3(PO4)2 lithium carbonate Li2CO3 2-D
Lecture Problem: II-8 (pg. 69) Naming A. Mg3(PO4)2 magnesium phosphate B. KHCO3 potassium bicarbonate Lecture Problem: II-9 (pg. 69) Formula Writing C. sodium carbonate Na2CO3 D. calcium sulfate CaSO4 2-D
Compounds with a variable oxidation state metal and a polyatomic anion Name the metal first, followed by the oxidation state (in Roman numerals) Then add the name of the polyatomic ion 2-D
Examples CuSO4 copper (II) sulfate Sn(NO3)4 tin (IV) nitrate nickel (III) bisulfate Ni(HSO4)3 titanium (IV) dichromate Ti(Cr2O7)2 2-D
Lecture Problem: II-10 (pg. 70) Naming Fe(C2H3O2)2 iron (II) acetate Lecture Problem: II-11 (pg. 70) Formula Writing tin (II) cyanide Sn(CN)2 2-D
Compounds having a polyatomic cation and monoatomic anion Name the positive polyatomic ion first, then the anion using the ide ending (NH4)3P ammonium phosphide ammonium sulfide (NH4)2S
Lecture Problem: II-12 (pg. 71) Naming NH4I ammonium iodide Lecture Problem: II-13 (pg. 71) Formula Writing ammonium oxide (NH4)2O 2-D
Polyatomic ions with other polyatomic ions Name the positive polyatomic ion first, then the negative polyatomic ion. Examples: ammonium carbonate ammonium nitrate (NH4)2CO3 NH4NO3
Lecture Problem: II-14 (pg. 71) Naming (NH4)2CrO4 ammonium chromate Lecture Problem: II-15 (pg. 71) Formula Writing ammonium oxide (NH4)2CO3 2-D
Naming Acids and Bases Acids are substances that release hydrogen ions in water. Binary acids Formed for a hydrogen cation and an ide anion Ternary acids Formed from a hydrogen cation and an ate anion
Binary Acids Binary acids are formed when an anion having the ide ending is paired with a hydrogen cation. To write the name use the hydro root name of the ide anion followed by ic acid Examples: HCl(aq) hydrochloric acid HF(aq) hydrofloric acid H2S(aq) hydrosulfuric acid
Lecture Problem: II-16 (pg. 73) Name these binary acids: HBr(aq) hydrobromic acid HI(aq) hydroiodic acid 2-D
Ternary Acids Consist of three different kinds of atoms Formed by making a compound from an -ate anion and a H+ cation To name the compound drop the ate and add ic acid Examples: H2SO4(aq) HNO3(aq) sulfuric acid (from sulfate) nitric acid (from nitrate)
Lecture Problem: II-17 (pg. 73) Name these ternary acids: H3PO4(aq) phosphoric acid H2CO3(aq) carbonic acid 2-D
Bases One definition of a base is that it produces OH-(aq) ions when dissolved in water. Bases are often formed by making a compound of a cation in combination with the OH- anion Bases react with acids to form water (neutralization) Bases are named just as with the polyatomic ions. (ternary compounds with hydroxide anion) Example: NaOH is sodium hydroxide
Part E: Chemical Bonding In review: Atoms are made up of a positive nucleus and negative electrons surrounding it. The chemical bonding properties of an element are largely determined by the electrons The periodic table is organized in order of increasing atomic number
Part E: Chemical Bonding In review: Atoms may either gain or loose electrons during bonding, forming ions Metals are the elements with the greatest tendency to loose electrons (cations) Nonmetals are the elements that tend to gain electrons (anions)
Part E: Chemical Bonding Electrons around an atom are arranged in what we call energy levels. The 1st energy level is closest to the nucleus As the energy levels increase they get further from the nucleus and closer together
Lewis Structures Lewis Structures provide a simple way to show the electrons in the outermost energy level. We don't write Lewis Structures for the transition elements, only the major groups (group A) elements.
Lewis Structures For a Group A element in its neutral state the number of valence electrons is equal to its group number. oxygen fluorine neon O F Ne
Ionic Bonds with Lewis Structures K F K + F - KF Let's practice drawing Lewis dot structures for the following ionic compounds. (Fill in on page 80) AlCl3 Na2O Ga2S3
Lewis Structures for Covalent Bonding In a covalent compound each atom attains stability by sharing electrons to fill its valence shell. Most atoms obey the octet rule: Stability is obtained when 8 electrons occupy an atoms valence shell
Lewis Structures for Covalent Bonding: The Method 1. Look at the Lewis structure for each atom in the molecule and add all valence electrons. Cl Cl 7+7 = 14 electrons
Lewis Structures for Covalent Bonding: The Method 2. Divide the number of valence electrons by 2, to get the number of electron pairs in the molecule. Cl Cl 7+7 = 14 electrons 14 valence electrons / 2 = 7 electron pairs
Lewis Structures for Covalent Bonding: The Method 3. Draw a single line between each atom to signify a pair of bonding electrons Cl Cl 4. Subtract the bonding pair(s) from the total electron pairs. 7 electron pairs 1 bonding pair = 6 electron pairs
Lewis Structures for Covalent Bonding: The Method 5. Arrange the remaining pairs around the bonded atoms, alternating between atoms Cl Cl 6. Now make sure each atom has an octect, and that you have placed all electrons.
Lewis Structures for Covalent Compounds Now lets try some slightly more complicated molecules... OF2 NH3 Now circle the bonding pairs of electrons and place a box around the lone pairs (lecture problem pg 84.)
What happens when some atoms can't fill their octet? Let's consider the compound sulfur dioxide. SO2 18 valence electrons = 9 electron pairs Sulfur will be our central atom O S O
What happens when some atoms can't fill their octet? O S O 9 electron pairs 2 bonding pairs = 7 electron pairs Now we place the remaining electron pairs...
What happens when some atoms can't fill their octet? O S O Do Nowallallthe atoms atoms have have a full a full octet octet???! NO! Sulfur only has 6 electrons in its valence shell Now let's try: CO2 O3 HCN SO3
Lewis Structures for polyatomic ions Sometimes multiple bonds will not provide each atom with a full octet. In some cases electrons will be given up or taken in order to fill the octets, creating ions. For Example: NO3 - OH - + NH4
Lewis Structures Lewis structures: Provide an easy way to visualize valence electrons Show us how atoms are bonded together in compounds Lewis structures do not tell us the actual molecular shape of a compound.