HCl = Hydrochloric Acid

HCl = Hydrochloric Acid
NaCl = Sodium Chloride Chapter 5 Nomenclature HCl = Hydrochloric Acid

Chemical BONDING

Chemical Bond A bond results from the attraction of nuclei for electrons All atoms trying to achieve a stable octet IN OTHER WORDS the p+ in one nucleus are attracted to the e- of another atom Electronegativity

Molecule: 2 or more atoms joined by a chemical bond
Compound: a molecule composed of atoms of 2 or more different elements bonded together in a fixed ratio

Diatomic Molecule Diatomic Molecule: a molecule containing 2 atoms
The Diatomic molecules are: Hydrogen (H2) Nitrogen (N2) Oxygen (O2) Fluorine (F2) Chlorine (Cl2) Iodine (I2) Bromine (Br2)

Chemical formula: represents the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numeric subscripts Bond energy: the energy required to break a chemical bond and form neutral atoms

Types of Chemical Bonds: (4)
Ionic bonds Covalent bonds Metallic bonds Hydrogen bonds Copyright © Cengage Learning. All rights reserved

What did the atom of fluorine
say to the atom of sodium? You complete me.

Bond Formation exothermic process ENERGY Reactants Energy released
Products

Breaking Bonds Endothermic reaction
energy must be put into the bond in order to break it ENERGY Products Energy Absorbed Reactants

Bond Strength Strong, STABLE bonds require lots of energy to be formed or broken weak bonds require little E

Is a bond forming or breaking? Strong or weak bond?
Reaction Time Energy (KJ) energy absorbed endothermic bond breaking weak unstable bond Products Reactants

Reaction Time Energy Energy absorbed endothermic bond breaking strong stable bond Products Reactants

Energy (KJ) Reaction Time Energy released exothermic bond formation weak unstable bond Reactants Products

Energy (KJ) Reaction Time Reactants Products Energy released exothermic bond formation strong stable bond

Two Major Types of Bonding
Ionic Bonding forms ionic compounds transfer of e- Covalent Bonding forms molecules sharing e-

One minor type of bonding
Metallic bonding Occurs between like atoms of a metal in the free state Valence e- are mobile (move freely among all metal atoms) Positive ions in a sea of electrons Metallic characteristics High mp temps, ductile, malleable, shiny Hard substances Good conductors of heat and electricity as (s) and (l)

It’s the mobile electrons that enable me-tals to conduct electricity!!!!!!

IONic Bonding electrons are transferred between valence shells of atoms ionic compounds are made of ions NOT MOLECULES ionic compounds are called Salts or Crystals

[METALS ]+ [NON-METALS ]-
IONic bonding Always formed between metals and non-metals [METALS ]+ [NON-METALS ]- Lost e- Gained e-

Anion (-) Cation (+)

Properties of Ionic Compounds
SALTS Crystals hard 22oC high mp temperatures nonconductors of electricity in solid phase good conductors in liquid phase or dissolved in water (aq)

Covalent Bonding Pairs of e- are shared between non-metal atoms
molecules Pairs of e- are shared between non-metal atoms electronegativity difference < 2.0 forms polyatomic ions

Properties of Molecular Substances
Covalent bonding Low m.p. temp and b.p. temps relatively soft solids as compared to ionic compounds nonconductors of electricity in any phase

Covalent, Ionic, metallic bonding?
NO2 sodium hydride Hg H2S sulfate NH4+ Aluminum phosphate KH KCl HF CO Co Can You Tell What type of bond is formed

Drawing ionic compounds using Lewis Dot Structures
Symbol represents the KERNEL of the atom (nucleus and inner e-) dots represent valence e-

[Na]+ [ Cl ]- NaCl This is the finished Lewis Dot Structure
How did we get here?

Step 1 after checking that it is IONIC
Determine which atom will be the +ion Determine which atom will be the - ion Step 2 Write the symbol for the + ion first. NO DOTS Draw the e- dot diagram for the – ion COMPLETE outer shell Step 3 Enclose both in brackets and show each charge

Draw the Lewis Diagrams
LiF MgO CaCl2 K2S

Drawing molecules using Lewis Dot Structures
Symbol represents the KERNEL of the atom (nucleus and inner e-) dots represent valence e-

Always remember atoms are trying to complete their outer shell!
The number of electrons the atoms needs is the total number of bonds they can make. Ex. … H? O? F? N? Cl? C? one two one three one four

Methane CH4 This is the finished Lewis dot structure
How did we get here?

Step 1 count total valence e- involved Step 2 connect the central atom (usually the first in the formula) to the others with single bonds Step 3 complete valence shells of outer atoms Step 4 add any extra e- to central atom IF the central atom has 8 valence e- surrounding it . . YOU’RE DONE!

Sometimes . . . You only have two atoms, so there is no central atom, but follow the same rules. Check & Share to make sure all the atoms are “happy”. Cl Br H O N HCl

O O N N DOUBLE bond TRIPLE bond atoms that share two e- pairs (4 e-)
atoms that share three e- pairs (6 e-) N N

Draw Lewis Dot Structures
You may represent valence electrons from different atoms with the following symbols x, , CO2 NH3

Draw the Lewis Dot Diagram for polyatomic ions
Count all valence e- needed for covalent bonding Add or subtract other electrons based on the charge REMEMBER! A positive charge means it LOST electrons!!!!!

Draw Polyatomics Ammonium Sulfate

Types of Covalent Bonds
NON-Polar bonds Electrons shared evenly in the bond E-neg difference is zero Between identical atoms Diatomic molecules

Types of Covalent Bonds
Polar bond Electrons unevenly shared

non-polar MOLECULES Sometimes the bonds within a molecule are polar and yet the molecule is non-polar because its shape is symmetrical. H C Draw Lewis dot first and see if equal on all sides

Polar molecules (a.k.a. Dipoles)
Not equal on all sides Polar bond between 2 atoms makes a polar molecule asymmetrical shape of molecule

H Cl + -

Water is asymmetrical + + H O -

Water is a bent molecule
H H

W - A - T - E - R as bent as it can be! Water’s polar MOLECULE!
The H is positive The O is not - not - not - not

Making sense of the polar non-polar thing
BONDS Non-polar Polar Identical Different MOLECULES Non-polar Polar Symmetrical Asymmetrical

IONIC bonds …. Ionic bonds are so polar that the electrons are not shared but transferred between atoms forming ions!!!!!!

4 Shapes of molecules

Linear (straight line)
Ball and stick model Space filling model

Bent Ball and stick model Space filling model

Trigonal pyramid Ball and stick model Space filling model

Tetrahedral Ball and stick model Space filling model

Intermolecular attractions
Attractions between molecules van der Waals forces Weak attractive forces between non-polar molecules Hydrogen “bonding” Strong attraction between special polar molecules

van der Waals Non-polar molecules can exist in liquid and solid phases
because van der Waals forces keep the molecules attracted to each other Exist between CO2, CH4, CCl4, CF4, diatomics and monoatomics

van der Waals periodicity
increase with molecular mass. Greater van der Waals force? F2 Cl2 Br2 I2 increase with closer distance between molecules Decreases when particles are farther away

Hydrogen “Bonding” Strong polar attraction
Like magnets Occurs ONLY between H of one molecule and N, O, F of another H “bond”

H is shared between 2 atoms of OXYGEN or 2 atoms of NITROGEN or 2 atoms of FLUORINE Of 2 different molecules

Why does H “bonding” occur?
Nitrogen, Oxygen and Fluorine small atoms with strong nuclear charges powerful atoms very high electronegativities

Intermolecular forces dictate chemical properties
Strong intermolecular forces cause high b.p., m.p. and slow evaporation (low vapor pressure) of a substance.

Which substance has the highest boiling point?
HF NH3 H2O WHY? Fluorine has the highest e-neg, SO HF will experience the strongest H bonding and  needs the most energy to weaken the i.m.f. and boil

Density????

H2O(s) is less dense than H2O(l)
The hydrogen bonding in water(l) molecules is random. The molecules are closely packed. The hydrogen bonding in water(s) molecules has a specific open lattice pattern. The molecules are farther apart.

Chemical Names and formulas
With all of the compounds and all of the elements to be identified, a systematic method for writing formulas and naming compounds is necessary A correctly written chemical formula must represent the known facts about the composition of a compound Care must be taken so that subscripts are correct Copyright © Cengage Learning. All rights reserved

Using Chemical formulas
Chemical formulas indicate the elements present in a compound and the relative numbers of atoms of each element in the compound In chemical formulas, the elements are given by their symbols and the relative number of atoms of each element by numerical subscript Ex H2SO4 the H, S & O are symbols, the 2 & 4 are subscripts Copyright © Cengage Learning. All rights reserved

Ion: A charged particle due to loss or gain of electrons
Cation: positive charge ion represented by a (+) after the chemical symbol (metal) Ex Na+ Anion: negative charge ion represented by a (-) after the chemical symbol (metal) Ex Cl-

Examples: F- fluoride ion S-2 sulfide ion I- Iodide ion
Negative ions are named by dropping the ending of the element name and adding the ending “ide” to it followed by the word “ion” Examples: F fluoride ion S sulfide ion I Iodide ion Copyright © Cengage Learning. All rights reserved

Learning Check Give the names of the following ions: Ba2+ Al3+ K+
_________ __________ _________ N3 O2 F _________ __________ _________ P3 S2 Cl

Solution Ba2+ Al3+ K+ barium aluminum potassium N3 O2 F
nitride oxide fluoride P3 S2 Cl phosphide sulfide chloride

Binary Ionic Compounds Binary Covalent Compounds
Binary Compounds Composed of two elements Binary Ionic Compounds Metal—nonmetal Binary Covalent Compounds Nonmetal—nonmetal Copyright © Cengage Learning. All rights reserved

Binary Ionic Compounds
Copyright © Cengage Learning. All rights reserved

Binary ionic compounds contain positive cations and negative anions.
Type I compounds Metal present forms only one cation. Type II compounds Metal present can form 2 or more cations with different charges. Copyright © Cengage Learning. All rights reserved

Metals (Groups I, II, and III) and Non-Metals
Type I Compounds Metals (Groups I, II, and III) and Non-Metals Metal _________ + Non-Metal _________ide Sodium Chlorine Sodium Chloride NaCl

Common Simple Cations and Anions

Rules for Naming Type I Ionic Compounds
1. The cation is always named first and the anion second. 2. A simple cation takes its name from the name of the element. A simple anion is named by taking the first part of the element name (the root) and adding –ide. Copyright © Cengage Learning. All rights reserved

Binary Ionic Compounds (Type I)
Examples: KCl Potassium chloride MgBr2 Magnesium bromide CaO Calcium oxide Copyright © Cengage Learning. All rights reserved

What is the name of the compound SrBr2? a) strontium bromine
Exercise What is the name of the compound SrBr2? a) strontium bromine b) sulfur bromide c) strontium dibromide d) strontium bromide Strontium bromide. Sr is the symbol for strontium. Br is the symbol for bromine, but take the first part of the element name (the root) and add –ide to get the name bromide. Copyright © Cengage Learning. All rights reserved

Strontium bromide. Sr is the symbol for strontium.
Br is the symbol for bromine, take the first part of the element name (the root) and add –ide to get the name bromide. Copyright © Cengage Learning. All rights reserved

Binary Ionic Compounds (Type II)
Metals in these compounds can form more than one type of positive charge. Charge on the metal ion must be specified. Roman numeral indicates the charge of the metal cation. Transition metal cations usually require a Roman numeral. Copyright © Cengage Learning. All rights reserved

Type II Compounds Metals (Transition Metals) and Non-Metals
Metal ______ +Roman Numeral (__) + Non-Metal ________ide Iron III Bromine Iron (III) Bromide FeBr3 Compare with Iron (II) Bromide FeBr2 Metals (Transition Metals) and Non-Metals Older System Metal (Latin) _______ + ous or ic + Non-Metal ________ide Ferrous Bromine Ferrous Bromide FeBr2 Compare with Ferric Bromide FeBr3

Old system: “ous” ending for lower charge
Different names are needed for positive ions of 2 different charges formed by the same metal Old system: “ous” ending for lower charge “ic” ending for higher charge New system: gives actual charge on the ion as a roman numeral Copyright © Cengage Learning. All rights reserved

Rules for Naming Type II Ionic Compounds
1. The cation is always named first and the anion second. 2. Because the cation can assume more than one charge, the charge is specified by a Roman numeral in parentheses. Copyright © Cengage Learning. All rights reserved

Binary Ionic Compounds (Type II)
Examples: CuBr Copper(I) bromide FeS Iron(II) sulfide PbO2 Lead(IV) oxide Copyright © Cengage Learning. All rights reserved

What is the name of the compound CrO2? a) chromium oxide
Exercise What is the name of the compound CrO2? a) chromium oxide b) chromium(II) oxide c) chromium(IV) oxide d) chromium dioxide Chromium(IV) oxide. Cr is the symbol for chromium. O is the symbol for oxygen, but take the first part of the element name (the root) and add –ide to get the name oxide. Since chromium can have more than one charge, a Roman numeral must be used to identify that charge. There are two oxygen ions each with a 2– charge, giving an overall charge of –4. Therefore, the charge on chromium must be +4. Copyright © Cengage Learning. All rights reserved

Therefore, the charge on chromium must be +4.
Chromium(IV) oxide. Cr is the symbol for chromium. O is the symbol for oxygen, but take the first part of the element name (the root) and add –ide to get the name oxide. Since chromium can have more than one charge, a Roman numeral must be used to identify that charge. There are two oxygen ions each with a 2– charge, giving an overall charge of –4. Therefore, the charge on chromium must be +4. Copyright © Cengage Learning. All rights reserved

b) chromium(II) sulfide c) nickel(III) sulfate d) iron(II) sulfide
Exercise What is the correct name of the compound that results from the most stable ion for sulfur and the metal ion that contains 24 electrons? a) iron(III) sulfide b) chromium(II) sulfide c) nickel(III) sulfate d) iron(II) sulfide Iron(II) sulfide. For sulfur, take the first part of the element name (the root) and add –ide to get the name sulfide. Iron with a +2 charge (as the Roman numeral indicates) contains 24 electrons (26p – 24e = +2 charge). Copyright © Cengage Learning. All rights reserved

Iron(II) sulfide. For sulfur, take the first part of the element name (the root) and add –ide to get the name sulfide. Iron with a +2 charge (as the Roman numeral indicates) contains 24 electrons (26p – 24e = +2 charge). Copyright © Cengage Learning. All rights reserved

Binary Covalent Compounds

Rules for Naming Type III Binary Compounds
Formed between two nonmetals. 1. The first element in the formula is named first, and the full element name is used. 2. The second element is named as though it were an anion. 3. Prefixes are used to denote the numbers of atoms present. 4. The prefix mono- is never used for naming the first element. Copyright © Cengage Learning. All rights reserved

Non-Metals and Non-Metals
Type III Compounds Non-Metals and Non-Metals Use Prefixes such as mono, di, tri, tetra, penta, hexa, hepta, etc. CO2 Carbon dioxide CO Carbon monoxide PCl3 Phosphorus trichloride CCl4 Carbon tetrachloride N2O5 Dinitrogen pentoxide CS2 Carbon disulfide Copyright © Cengage Learning. All rights reserved

Prefixes Used to Indicate Numbers in Chemical Names
Additional Prefixes 9 nona- 10 deca- 11 undeca- 12 dodeca- 13 trideca- 14 tetradeca- 15 pentadeca- 16 hexadeca- 17 heptadeca- 18 octadeca- 19 nonadeca- 20 icosa Copyright © Cengage Learning. All rights reserved

Binary Covalent Compounds (Type III)
Examples: CO2 Carbon dioxide SF6 Sulfur hexafluoride N2O4 Dinitrogen tetroxide Copyright © Cengage Learning. All rights reserved

What is the name of the compound SeO2? a) selenium oxide
Exercise What is the name of the compound SeO2? a) selenium oxide b) selenium dioxide c) selenium(II) oxide d) selenium(IV) dioxide Selenium dioxide. Se is the symbol for selenium. O is the symbol for oxygen, but take the first part of the element name (the root) and add –ide to get the name oxide. Since they are both nonmetals, prefixes are used to identify the elements (except mono- is not used for the first element). Two oxygen atoms require the use of the prefix di-, making the name dioxide. Copyright © Cengage Learning. All rights reserved

Se is the symbol for selenium. O is the symbol for oxygen,
Selenium dioxide. Se is the symbol for selenium. O is the symbol for oxygen, take the first part of the element name (the root) and add –ide to get the name oxide. Since they are both nonmetals, prefixes are used to identify the elements (except mono- is not used for the first element). Two oxygen atoms require the use of the prefix di-, making the name dioxide. Copyright © Cengage Learning. All rights reserved

Flow Chart for Naming Binary Compounds

Let’s Practice! Name the following. CaF2 Calcium Flouride K2S Potassium Sulfide CoI2 Cobalt (II) Iodide or Cobaltous Iodide SnF2 Tin (II) Fluoride or Stannous Fluoride SnF4 Tin (IV) Fluoride or Stannic Fluoride OF2 Oxygen diflouride CuI2 Copper (II) Iodide or Cupric Iodide CuI Copper (I) Iodide or Cuprous Iodide SO2 Sulfur dioxide SrS Strontium Sulfide Lithium Bromide LiBr

Naming Polyatomic compounds

They have special names and must be memorized.
Polyatomic ions are charged entities composed of several atoms bound together. They have special names and must be memorized. We will be using our Fat Daddy Chart to help us with naming the polyatomic compounds Those used often enough will be memorized just out of sheer practice Copyright © Cengage Learning. All rights reserved

Names of Common Polyatomic Ions (page 130)

Mg(NO3)2 Magnesium nitrate (NH4)2SO4 Ammonium sulfate
Examples NaOH Sodium hydroxide Mg(NO3)2 Magnesium nitrate (NH4)2SO4 Ammonium sulfate FePO4 Iron(III) phosphate Copyright © Cengage Learning. All rights reserved

Learning Check Select the correct name for each. A. Fe2S3
1) iron sulfide 2) iron(II) sulfide 3) iron(III) sulfide B. CuO 1) copper oxide 2) copper(I) oxide 3) copper(II) oxide

Solution Select the correct name for each. A. Fe2S3
3) iron(III) sulfide Fe3+ S2– B. CuO 3) copper(II) oxide Cu2+ O2–

Overall Strategy for Naming Chemical Compounds

What is the name of the compound KClO3? a) potassium chlorite
Exercise What is the name of the compound KClO3? a) potassium chlorite b) potassium chlorate c) potassium perchlorate d) potassium carbonate ClO3– is the polyatomic ion chlorate. Copyright © Cengage Learning. All rights reserved

Exercise Examine the following table of formulas and names. Which of the compounds are named correctly? a) I, II b) I, III, IV c) I, IV d) I only Formula Name I P2O5 Diphosphorus pentoxide II ClO2 Chlorine oxide III PbI4 Lead iodide IV CuSO4 Copper(I) sulfate Only Formula I is named correctly. Formula II is chlorine dioxide. Formula III is lead(IV) iodide. Formula IV is copper(II) sulfate. Copyright © Cengage Learning. All rights reserved

Only Formula I is named correctly. Formula II is chlorine dioxide.
Formula III is lead(IV) iodide. Formula IV is copper(II) sulfate. Copyright © Cengage Learning. All rights reserved

Molecule with one or more H+ ions attached to an anion.
Acids Acids can be recognized by the hydrogen that appears first in the formula—HCl. Molecule with one or more H+ ions attached to an anion. Most lab acids are either: binary acids ( composed of Hydrogen and another element) or oxyacids (composed of Hydrogen, oxygen and a third element Copyright © Cengage Learning. All rights reserved

Rules for Naming Acids If the anion does not contain oxygen, the acid is named with the prefix hydro– and the suffix –ic attached to the root name for the element. Examples: HCl Hydrochloric acid HCN Hydrocyanic acid H2S Hydrosulfuric acid Copyright © Cengage Learning. All rights reserved

Acids That Do Not Contain Oxygen

If the anion contains oxygen:
Rules for Naming Acids If the anion contains oxygen: The suffix –ic is added to the root name if the anion name ends in –ate. Examples: HNO3 Nitric acid H2SO4 Sulfuric acid HC2H3O2 Acetic acid Copyright © Cengage Learning. All rights reserved

If the anion contains oxygen:
Rules for Naming Acids If the anion contains oxygen: The suffix –ous is added to the root name if the anion name ends in –ite. Examples: HNO2 Nitrous acid H2SO3 Sulfurous acid HClO2 Chlorous acid Copyright © Cengage Learning. All rights reserved

Some Oxygen-Containing Acids

Flowchart for Naming Acids

Which of the following compounds is named incorrectly?
Exercise Which of the following compounds is named incorrectly? KNO3 potassium nitrate TiO2 titanium(II) oxide Sn(OH)4 tin(IV) hydroxide PBr5 phosphorus pentabromide H2SO3 sulfurous acid The correct answer is “b”. The charge on oxygen is 2–. Since there are two oxygen atoms, the overall charge is 4–. Therefore, the charge on titanium must be 4+ (not 2+ as the Roman numeral indicates). Copyright © Cengage Learning. All rights reserved

The correct answer is “b”. The charge on oxygen is 2–.
Since there are two oxygen atoms, the overall charge is 4–. Therefore, the charge on titanium must be 4+ (not 2+ as the Roman numeral indicates). Copyright © Cengage Learning. All rights reserved

Examples Sodium hydroxide NaOH Potassium carbonate K2CO3 Sulfuric acid
H2SO4 Dinitrogen pentoxide N2O5 Cobalt(III) nitrate Co(NO3)3 Copyright © Cengage Learning. All rights reserved

a) phosphorus trichloride b) carbon monochloride c) tin(IV) chloride
Exercise A compound has the formula XCl3 where X could represent a metal or nonmetal. What could the name of this compound be? a) phosphorus trichloride b) carbon monochloride c) tin(IV) chloride d) magnesium chloride Phosphorus trichloride. Carbon monochloride has the formula CCl. Tin(IV) chloride has the formula SnCl4. Magnesium chloride has the formula MgCl2. Phosphorus trichloride has the formula PCl3 and is therefore the correct answer. Copyright © Cengage Learning. All rights reserved

Phosphorus trichloride. Carbon monochloride has the formula CCl.
Tin(IV) chloride has the formula SnCl4. Magnesium chloride has the formula MgCl2. Phosphorus trichloride has the formula PCl3 and is therefore the correct answer Copyright © Cengage Learning. All rights reserved

Lets Practice Some More!
HF Hydroflouric acid Na2CO3 Sodium carbonate H2CO3 Carbonic acid KMnO4 Potassium permanganate HClO4 Perchloric acid H2S Hyrdogen sulfuric acid NaOH Sodium hydroxide CuSO4 Copper (II) sulfate or Cupric sulfate PbCrO4 Lead (II) chromate or Plubous chromate H2O Hydrooxic acid (no……just water) NH3 Nitrogen trihydride (no..just ammonia)

Writing Chemical formulas

Identifying Ionic Charges
Group A elements – use the periodic table to determine ionic charge * elements in same group have same ionic charge * Group 4A and Noble gases – almost never form ions Group B elements – many have more than one ionic charge Periodic Table For Group A elements, the periodic table can be used to determine whether an element will form positive cations or negative anions. The P.T. also can be used to determine charge on these ions Elements in the same group have the same charge when they form ions. Group 4A and the Noble gases are the exception. They almost never form ions. For the Group B elements (transition metals), many have more than one common ionic charge.

Identifying Ionic Charges
Periodic Table For Group A elements, the periodic table can be used to determine whether an element will form positive cations or negative anions. The P.T. also can be used to determine charge on these ions Elements in the same group have the same charge when they form ions. Group 4A and the Noble gases are the exception. They almost never form ions. For the Group B elements (transition metals), many have more than one common ionic charge. The alkali metals reliably form +1 ions; alkaline earth metals form +2 ions; Group 6 elements form -2 ions, and halogens form -1 ions. Elements in Group 3 usually form +3 ions, and other elements vary in their ion charges. Metals generally form positive ions, and nonmetals form negative ions. Main group elements tend to form only one type of ion, while the transition metals tend to form two or more types of ions. Charge on cations corresponds to group #. Charge on anions is found by subtracting 8 by group number the number 8 is used b/c it represents # of valence e- in Noble gases

Naming Cations and Anions
Stop Naming Cations and Anions Write the name or symbol w/ charge: Potassium ion Copper (II) ion Chloride ion Oxide ion Ba2+ S2- Au3+ Nitrite ion Hydroxide ion Phosphate ion SO42- CrO42- ClO32- Practice writing names and formulas of ions by writing the answers in your notebook.

Binary Ionic Compounds
Compounds composed of 2 different monatomic elements To write binary formulas – write cation first, then anion *criss-cross charges to determine how many of each ion you need *use subscripts to denote number of ions ex: Ca2+ + Cl CaCl2 Na1+ + Cl NaCl Binary ionic compounds – are compounds composed of two different monatomic elements To write formulas for binary ionic compounds, write the symbol for the cation first, followed by the anion. Cross-over the charges by using the absolute value of each ion’s charge as a subscript for the other ion. If there is only one ion needed in a formula, a subscript is not needed (don’t write “1”) To name binary ionic compounds, write the name of the cation first, followed by the name of the anion with the –ide ending

Ternary Ionic Compounds
Compounds containing at least one polyatomic ion; at least 3 different elements To write ternary formulas: write cation first, then anion *criss-cross charges to determine how many of each ion you need *use subscripts to denote number of ions *must use parentheses around polyatomic if more than one is needed!!! ex: Na1+ + SO Na2SO3 Mg2+ + OH Mg(OH)2 [not same as MgOH2] Ternary ionic compounds – compounds containing at least one polyatomic ion; have at least 3 different elements To write the formulas for ternary ionic compounds, use the same rules as for binary compounds Except: use parentheses when more than one polyatomic ion is needed; use criss-cross method like binary To name ternary ionic compounds, use the name rules as for binary compounds except remember that some polyatomic anions do not end in –ide. Be sure to note the exceptions.

Write the name or the formula for the following compounds:
Stop Ionic Compounds Write the name or the formula for the following compounds: NaNO3 CaSO4 (NH4)2O CuSO3 Fe(OH)3 NaF Lithium sulfide Iron (III) phosphide Magnesium fluoride Barium nitrate Aluminum hydroxide Potassium phosphate Practice writing names and formulas of ionic compounds by writing the answers in your notebook. Practice making ionic compounds!

Write the name or formula for the following compounds:
Stop Molecular Compounds Write the name or formula for the following compounds: P2O5 N2O NO2 CBr4 CO2 tetraiodine nonoxide sulfur hexafluoride nitrogen trioxide carbon tetrahydride phosphorus trifluoride Now practice writing the correct name or formula for the following molecular compounds in your notebook.

Examples of Ionic Compounds with Two Elements
Formula Ions Name Cation Anion NaCl Na+ Cl– sodium chloride K2S K S2– potassium sulfide MgO Mg2+ O2– magnesium oxide CaI2 Ca2+ I– calcium iodide Al2O3 Al S2– aluminum sulfide

Learning Check Br– S2− N3− Na+ Al3+
Write the formulas and names for compounds of the following ions: Br– S2− N3− Na+ Al3+

Solution Na+ Al3+ NaBr sodium bromide Na2S sodium sulfide Na3N
Br− S2− N3− Na+ Al3+ NaBr sodium bromide Na2S sodium sulfide Na3N sodium nitride AlBr3 aluminum bromide Al2S3 aluminum sulfide AlN aluminum nitride

Transition Metals Form Positive Ions
Most transition metals and Group 4(14) metals, Form 2 or more positive ions Zn2+, Ag+, and Cd2+ form only one ion.

Guide to Writing Formulas from the Name

Writing Formulas Write a formula for potassium sulfide.
STEP 1 Identify the cation and anion. potassium = K+ sulfide = S2− STEP 2 Balance the charges. K S2− K+ 2(1+) + 1(2–) = 0 STEP 3 Write the cation first. 2K+ and 1S2− = K2S1 = K2S

Writing Formulas Write a formula for iron(III) chloride.
STEP 1 Identify the cation and anion. iron (III) = Fe3+ (III = charge of 3+) chloride = Cl− STEP 2 Balance the charges. Fe Cl− Cl− 1(3+) + 3(1–) = 0 STEP 3 Write the cation first. 1Fe3+ and 3Cl− = FeCl3

Learning Check The correct formula for each of the following is:
A. copper(I) nitride 1) CuN 2) CuN3 3) Cu3N B. lead(IV) oxide 1) PbO2 2) PbO 3) Pb2O4

Solution The correct formula for each of the following is:
A. copper(I) nitride 3) Cu3N 3Cu+ + N3– = 3(1+) + (3–) = 0 B. lead(IV) oxide 1) PbO2 Pb O2– = (4+) + 2(2–) = 0

Percent Composition, Empirical Formulas, Molecular Formulas

Formula Masses and Molar masses:
Molecular mass or molecular weight are used instead of the term formula mass. The formula mass of any compound is the sum of the average atomic masses of all of the atoms present in the formula

Example of formula mass
H2O 2 H atom weigh each 1 O atom weighs each 2 x 1.oo79 +1x formula mass for water

Molar mass as a conversion factor
Moles x grams/mole = mass in grams Mass in grams x 1 mol/grams = moles Thus 2 conversions relate mass in grams to numbers of moles of a substance Copyright © Cengage Learning. All rights reserved

What is the molar mass of Barium nitrate Ba(NO3)2
Example What is the molar mass of Barium nitrate Ba(NO3)2 Solution 1 mol Ba x g/1 mol Ba = g Ba 2 moles N x g/1mole N = g N 6 moles O x g/1mol O = g Molar mass Ba(NO3)2 = Copyright © Cengage Learning. All rights reserved

So… Percent Composition Part _______ Percent = x 100% Whole
Percent Composition – the percentage by mass of each element in a compound Part _______ Percent = x 100% Whole So… Percent composition of a compound or = molecule Mass of element in 1 mol ____________________ x 100% Mass of 1 mol

Percent Composition Molar Mass of KMnO4 K = 1(39.1) = 39.1
Example: What is the percent composition of Potassium Permanganate (KMnO4)? Molar Mass of KMnO4 K = 1(39.1) = 39.1 Mn = 1(54.9) = 54.9 O = 4(16.0) = 64.0 MM = 158 g

Percent Composition Molar Mass of KMnO4 = 158 g 39.1 g K % K x 100 =
Example: What is the percent composition of Potassium Permanganate (KMnO4)? Molar Mass of KMnO4 = 158 g 39.1 g K % K x 100 = 24.7 % 158 g 54.9 g Mn 34.8 % x 100 = % Mn 158 g K = 1(39.10) = 39.1 64.0 g O x 100 = 40.5 % Mn = 1(54.94) = 54.9 % O 158 g O = 4(16.00) = 64.0 MM = 158

Percent Composition Determine the percentage composition of sodium carbonate (Na2CO3)? Molar Mass Percent Composition 46.0 g x 100% = 43.4 % Na = 2(23.00) = 46.0 C = 1(12.01) = 12.0 O = 3(16.00) = 48.0 MM= g % Na = 106 g 12.0 g x 100% = 11.3 % % C = 106 g 48.0 g x 100% = 45.3 % % O = 106 g

Percent Composition Determine the percentage composition of ethanol (C2H5OH)? % C = 52.13%, % H = 13.15%, % O = 34.72% _______________________________________________ Determine the percentage composition of sodium oxalate (Na2C2O4)? % Na = 34.31%, % C = 17.93%, % O = 47.76%

Percent Composition Calculate the mass of bromine in 50.0 g of Potassium bromide. 1. Molar Mass of KBr K = 1(39.10) = 39.10 Br =1(79.90) =79.90 MM = 119.0 2. 79.90 g ___________ = 119.0 g x 50.0g = 33.6 g Br

Percent Composition Calculate the mass of nitrogen in 85.0 mg of the amino acid lysine, C6H14N2O2. 1. Molar Mass of C6H14N2O2 C = 6(12.01) = 72.06 H =14(1.01) = 14.14 N = 2(14.01) = 28.02 O = 2(16.00) = 32.00 MM = 146.2 2. 28.02 g ___________ = 0.192 146.2 g x 85.0 mg = 16.3 mg N

Hydrates Hydrated salt – salt that has water molecules trapped within the crystal lattice Examples: CuSO4•5H2O , CuCl2•2H2O Anhydrous salt – salt without water molecules Examples: CuCl2 Can calculate the percentage of water in a hydrated salt.

Percent Composition Calculate the percentage of water in sodium carbonate decahydrate, Na2CO3•10H2O. 1. Molar Mass of Na2CO3•10H2O Na = 2(22.99) = 45.98 C = 1(12.01) = 12.01 H = 20(1.01) = 20.2 3. O = 13(16.00)= 180.2 g MM = 286.2 _______ x 100%= 67.97 % 286.2 g 2. Water H = 20(1.01) = 20.2 O = 10(16.00)= MM = 180.2 H = 2(1.01) = 2.02 or So… 10 H2O = 10(18.02) = 180.2 O = 1(16.00) = 16.00 MM H2O = 18.02

Percent Composition Calculate the percentage of water in Aluminum bromide hexahydrate, AlBr3•6H2O. 1. Molar Mass of AlBr3•6H2O Al = 1(26.98) = 26.98 Br = 3(79.90) = 239.7 H = 12(1.01) = 12.12 O = 6(16.00) = 96.00 MM = 374.8 3. 108.1 g 2. Water _______ x 100%= 28.85 % 374.8 g H = 12(1.01) = 12.1 O = 6(16.00)= 96.00 MM = 108.1 MM = 18.02 For 6 H2O = 6(18.02) = 108.2 or

Percent Composition If 125 grams of magnesium sulfate heptahydrate is completely dehydrated, how many grams of anhydrous magnesium sulfate will remain? MgSO4 . 7 H2O 1. Molar Mass 2. % MgSO4 Mg = 1 x = g S = 1 x = g O = 4 x = g MM = g 120.4 g X 100 = 48.84 % 246.5 g 3. Grams anhydrous MgSO4 H = 2 x = 2.02 g O = 1 x = g MM = g x 125 = 61.1 g MM H2O = 7 x g = g Total MM = 120.4 g g = g

Percent Composition If 145 grams of copper (II) sulfate pentahydrate is completely dehydrated, how many grams of anhydrous copper sulfate will remain? CuSO4 . 5 H2O 1. Molar Mass 2. % CuSO4 Cu = 1 x = g S = 1 x = g O = 4 x = g MM = g 159.6 g X 100 = 63.92 % 249.7 g 3. Grams anhydrous CuSO4 H = 2 x = 2.02 g O = 1 x = g MM = g x 145 = 92.7 g MM H2O = 5 x g = 90.1 g Total MM = 159.6 g g = g

Percent Composition A 5.0 gram sample of a hydrate of BaCl2 was heated, and only 4.3 grams of the anhydrous salt remained. What percentage of water was in the hydrate? 2. Percent of water 1. Amount water lost 0.7 g water 5.0 g hydrate 4.3 g anhydrous salt 0.7 g water x 100 = 14 % 5.0 g hydrate

Percent Composition A 7.5 gram sample of a hydrate of CuCl2 was heated, and only 5.3 grams of the anhydrous salt remained. What percentage of water was in the hydrate? 2. Percent of water 1. Amount water lost 2.2 g water 7.5 g hydrate 5.3 g anhydrous salt 2.2 g water x 100 = 29 % 7.5 g hydrate

Percent Composition A 5.0 gram sample of Cu(NO3)2•nH2O is heated, and 3.9 g of the anhydrous salt remains. What is the value of n? 1. Amount water lost 5.0 g hydrate 3.9 g anhydrous salt 1.1 g water 3. Amount of water 0.22 x = 4.0 2. Percent of water 1.1 g water x 100 = 22 % 5.0 g hydrate

Percent Composition A 7.5 gram sample of CuSO4•nH2O is heated, and 5.4 g of the anhydrous salt remains. What is the value of n? 1. Amount water lost 7.5 g hydrate 5.4 g anhydrous salt 2.1 g water 3. Amount of water 0.28 x = 5.0 2. Percent of water 2.1 g water x 100 = 28 % 7.5 g hydrate

Formulas Percent composition allow you to calculate the simplest ratio among the atoms found in compound. Empirical Formula – formula of a compound that expresses lowest whole number ratio of atoms. Molecular Formula – actual formula of a compound showing the number of atoms present Examples: C6H12O6 - molecular C4H10 - molecular C2H5 - empirical CH2O - empirical

Formulas Is H2O2 an empirical or molecular formula?
Molecular, it can be reduced to HO HO = empirical formula

Calculating Empirical Formula
An oxide of aluminum is formed by the reaction of g of aluminum with g of oxygen. Calculate the empirical formula. 1. Determine the number of grams of each element in the compound. 4.151 g Al and g O 2. Convert masses to moles. 1 mol Al 4.151 g Al = mol Al 26.98 g Al 1 mol O 3.692 g O = mol O 16.00 g O

An oxide of aluminum is formed by the reaction of g of aluminum with g of oxygen. Calculate the empirical formula. 3. Find ratio by dividing each element by smallest amount of moles. moles Al = mol Al 0.1539 moles O = mol O 0.1539 4. Multiply by common factor to get whole number. (cannot have fractions of atoms in compounds) O = x 2 = 3 Al = x 2 = 2 therefore, Al2O3

A g sample of cobalt reacts with g chlorine to form a binary compound. Determine the empirical formula for this compound. 4.550 g Co 1 mol Co = mol Co 58.93 g Co 1 mol Cl 5.475 g Cl = mol Cl 35.45 g Cl mol Co mol Cl = 1 = 2 CoCl2

When a g sample of iron metal is heated in air, it reacts with oxygen to achieve a final mass of g. Determine the empirical formula. Fe = g O = g – g = g 1 mol Fe 2.000 g Fe = mol Fe 55.85 g Fe 1 mol O 0.573 g O = mol Fe 16.00 g 1 : 1 FeO 176

A sample of lead arsenate, an insecticide used against the potato beetle, contains g lead, g of hydrogen, g of arsenic, and g of oxygen. Calculate the empirical formula for lead arsenate. 1 mol Pb g Pb = mol Pb 207.2 g Pb gH 1 mol H = mol H 1.008 g H 1 mol As g As = mol As 74.92 g As 1 mol O 0.4267g Fe = mol O 16.00 g O

A sample of lead arsenate, an insecticide used against the potato beetle, contains g lead, g of hydrogen, g of arsenic, and g of oxygen. Calculate the empirical formula for lead arsenate. mol Pb = mol Pb mol H = 1.00 mol H PbHAsO4 mol As = mol As mol O = mol O

The most common form of nylon (Nylon-6) is 63.38% carbon, 12.38% nitrogen, 9.80% hydrogen and 14.14% oxygen. Calculate the empirical formula for Nylon-6. Step 1: In g of Nylon-6 the masses of elements present are g C, g n, 9.80 g H, and g O. Step 2: 63.38 g C 1 mol C 9.80 g H 1 mol H = mol C = 9.72 mol H 12.01 g C 1.01 g H 12.38 g N 1 mol N = mol N 14.01 g N 14.14 g O 1 mol O = mol O 16.00 g O

The most common form of nylon (Nylon-6) is 63.38% carbon, 12.38% nitrogen, 9.80% hydrogen and 14.14% oxygen. Calculate the empirical formula for Nylon-6. Step 3: 5.302 mol C = mol C 6:1:11:1 0.8837 mol N = mol N 0.8837 C6NH11O 9.72 mol H = 11.0 mol H 0.8837 mol O = mol O 0.8837

Calculating molecular formula
It is not possible to determine the correct molecular formula unless the molecular mass of the substance has been determined The relationship between the simplest formula and the molecular mass is: (simple formula)x = molecular formula Where x is a whole number multiple of the simple formula

Calculating Molecular Formula
A white powder is analyzed and found to have an empirical formula of P2O5. The compound has a molar mass of g. What is the compound’s molecular formula? Step 3: Multiply Step 1: Molar Mass P = 2 x g = 61.94g O = 5 x 16.00g = g g (P2O5)2 = P4O10 Step 2: Divide MM by Empirical Formula Mass g = 2 141.94g

Calculating Molecular Formula
A compound has an experimental molar mass of 78 g/mol. Its empirical formula is CH. What is its molecular formula? (CH)6 = C = g H = g 13.01 g C6H6 78 g/mol = 6 13.01 g/mol

Oxidation Numbers Are used to indicate general distributions of electrons among bonded atoms. Refer to handout for rules of oxidation numbers

Ex find oxidation # of following:
UF6 ClO3- Solution U = F = -1 Cl = O =-2

HCl = Hydrochloric Acid

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