Chapter 2: Compounds and Chemical Reactions

1. Chapter 2: Compounds and Chemical Reactions • Essentially all elements combine to form compounds • Compounds are of two types: • Molecular, which involve shared electrons and consist of electrically neutral, discrete particles called molecules • Ionic compounds, which involve electron transfer and charged particles called ions

2. Chemical formulas are collections of chemical symbols that are used to describe elements and compounds • Free elements are not combined with other elements in a compound • Examples: Fe (iron), Na (sodium), and K (potassium) • Many nonmetals occur as diatomic molecules

3. Chemical formulas specify the composition of a substance • NaCl is composed of the elements sodium and chlorine in a one-to-one (atom) ratio • Fe2O3 is composed of the elements iron and oxygen in a two-to-three ratio • CO(NH2)2 expands to CON2H4, but there are good reasons to write some compounds with parentheses • Hydrates are crystals that contain water molecules, for example plaster: CaSO4•2H2O • When all the water is removed (by heating), the solid that remains is said to be anhydrous (without water)

4. Chemical equations describe what happens in a chemical reactions • Hydrogen and oxygen combine to form water • Hydrogen and oxygen are called reactants • Water is called the product • Reactants are separated from products with “” 2 H2 + O2 2 H2O • Note that the “” is like an equal sign because both sides of the equation have the same number of each type of atom

5. Note: Mass is conserved because the number of atoms of each type remains the same on each side of the arrow. (Both sides of the arrow show 4 H and 2 O atoms.) This equation is said to be balanced. • This can be represented as: The “2” in front of formulas H2 and H2O are called coefficients. They indicate the number of molecules of each type and can change when balancing a chemical equation. The “2” in the formulas H2 and H2O indicate atom ratios for the compound and must not change.

6. It is sometimes useful to include the physical state of reactants and products • For solids use s, liquids use l, gases use g, and for aqueous solutions use aq. • For example, the reaction between stomach acid (an aqueous solution of HCl) and sodium carbonate (an antacid) can be written 2 HCl(aq) + CaCO3(s)  CaCl2(aq) + H2O(l) +CO2(g) Note: You can verify this equation is balanced by checking for mass balance: each side shows 1 Ca, 1 C, 2 Cl, 2 H, and 3 O.

7. Almost all chemical reactions either absorb or give off energy, often as heat or light • Kinetic and potential energy are both important in chemistry • Kinetic energy is the energy an object has when moving • Potential energy is the energy an object has due to its position • Potential energy is “stored energy” because it can be converted into kinetic energy

8. Energy must also be conserved • The Law of Conservation of Energy: • Energy cannot be created or destroyed; it can only be converted from one form to another • Heat and temperature are related to kinetic energy • The temperature of an object is proportional to its average kinetic energy (average speed of its atoms) • Heat or thermal energy is transferred between objects with different temperatures • Heat flow spontaneously from hot to cold objects

9. Chemical energy is a form of potential energy • The analysis of temperature changes in chemical reactions can provide information about the potential energy changes that occur • The kinetic molecular theory of matter provides more details about chemical energy changes and is discussed in Chapter 7 • Energy can also be transferred as light, which will be covered later in the book

10. As a general rule, molecular compounds are formed when nonmetallic elements combine • Many molecular compounds contain hydrogen: Group Noble Period IVA VA VIA VIIA Gas 2 CH4 NH3 H2O HF Ne 3 SiH4 PH3 H2S HCl Ar 4 GeH4 AsH3 H2Se HBr Kr 5 SbH3 H2Te HI Xe Note: The number of hydrogens that combined with the nonmetal equals the number of spaces to the right we have to move to get to the noble gas

11. Organic chemistry is a major specialty that deals with compounds containing mostly carbon and hydrogen • Hydrocarbons contain only hydrogen and carbon and are organic compounds • Alkanes are the simplest hydrocarbons • General formula is CnH2n+2 Space-filling models: Black atoms = carbon White atoms = hydrogen

12. Other classes of hydrocarbons exist • Different classes of organic compounds are derived from hydrocarbons by replacing hydrogen • For example alcohols result when a H is replaced by OH in a hydrocarbon Methanol (wood alcohol), CH3OH, is related to methane, CH4, by replacing one H with OH.

13. Inorganic compounds are substances not considered to be derived from hydrocarbons • The rules for naming, or nomenclature, of simple inorganic compound is covered now (organic nomenclature is covered later) • Binary compounds are compounds comprised of two different elements • The goal is to be able to convert between the chemical formula and the name

14. The first element in the formula is identified by its English name, the second by appending the suffix –ide to its stem Chemical Name as Name as Symbol Stem First Element Second Element O ox- oxygen oxide S sulf- sulfur sulfide N nitr- nitrogen nitride P phosph- phosphorus phosphide F fluor- fluorine fluoride Cl chlor- chlorine chloride Br brom- bromine bromide I iod- iodine iodide

15. Greek Prefixes mono- = 1 (often omitted) hexa- = 6 di- = 2 hepta- = 7 tri- = 3 octa- = 8 tetra- = 4 nona- = 9 penta- = 5 deca- = 10 • The number of each type of atom is specified with Greek prefixes Examples: PF5 = phosphorus pentafluoride HCl = hydrogen chloride N2O5 = dinitrogen tetraoxide Note: many compounds have common names, like water for H2O.

16. The subscripts in the formula of an ionic compound always specifies the smallest whole-number ratio of the ions because molecules don’t exist in ionic compounds • The smallest unit of a compound is called the formula unit • Positively charged ions have more protons than electrons and are called cations • Negatively charged ions have more electrons than protons and are called anions • The formula unit of an ionic compound always contains both cations and anions

17. Ionic compounds are composed of charged particles (ions) • Ions can be formed from the reaction of metal with a nonmetal • The metals form cations and the nonmetals form anions

18. The charges on many representative elements can be predicted: • Metals form cations • The positive charge on the cation is the same as the “A” group number of the metal • Nonmetals form anions • The negative charge on the anion is equal to the number of spaces to the right we have to move in the periodic table to get to a noble • Ionic compounds must be electrically neutral

19. Rules for writing Formulas of Ionic Compounds: 1) The positive ion is given first in the formula. 2) The subscripts in the formula must produce an electrically neutral formula unit. 3) The subscripts should be the set of smallest whole numbers possible. 4) The charges on the ions are not included in the finished formula of the substance.

20. Ions formed by transition metals (Group IIIB – VIIIB) and post-transition metals:

21. Some polyatomic ions (ions with two or more atoms): See Table 2.5 for a more polyatomic ions

22. Naming ionic compounds • The name of the cation is given first followed by the name of the anion • Cations: • If the metal forms only one positive ion, the cation name is the English name for the metal • If the metal forms more than one positive ion, the cation name is the English name followed, without a space, by the numerical value of the charge written as a Roman numeral in parentheses (this is for the Stock system) • Anions: • For monoatomic anions, the name is created by adding the “–ide” suffix to the stem name for the element. • For polyatomic ions, use the names in Table 2.5

23. To name a compound, you can use this flowchart:

24. Summary of Properties • Hardness and brittleness • Molecular compounds tend to be soft and easily crushed because the attractions between molecules are weak and molecules can slide past each other • Ionic compounds are hard and brittle because of the strong attractions and repulsions between ions

25. Melting points • To melt the particles in the solid must have sufficient kinetic energy to overcome the attractions between particles • Molecular compounds tend to have weak attractions between particles and so tend to have low melting points • Many molecular compounds are gases at room temperature • Ionic compound tend to have strong attractions so they have high melting points • Nearly all ionic compounds are solids at room temperature

26. Electrical conductivity requires the movement of electrical charge • Ionic compounds: • Do not conduct electricity in the solid state • Do conduct electricity in the liquid state • The ions are free to move in the liquid state • Molecular compounds: • Do not conduct electricity in the solid or liquid state • Molecules are comprised of uncharged particles