PassGAMSAT Science EClass 2

PassGAMSAT Science eClass 2

Chemical Bonding and Reactions Chemistry is the branch of science that deals with the composition, structure and properties of substances and how they change and interact with other substances and the environment. Physical chemistry explores how substances interact at the atomic level. Learning Objectives Students should be able to define, differentiate between and do calculations from the following topics: 1. Chemical bonds ionic covalent nomenclature for compounds 2. Chemical formulas formula weight 3. Chemical equations the mole calculations involving molarity writing and balancing equations 4. Solutions and solubility molarity 5. Types of reactions Redox reactions and oxidation number CHEMICAL BONDS A chemical bond is the attractive force which holds two or more atoms within a molecule. Within a given compound, each molecule is exactly the same as all of the others; the bonded atoms are the same distance apart (bond length) and arranged with the same geometry (bond angle).

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The amount of energy required to break a bond is called the bonding energy. All chemical bonds are the result of electrical attraction between atoms. They involve atoms either "giving", "receiving", or sharing electrons in order to reach the most stable noble gas electron configuration, which is to have the outermost or valence shell either completely filled or completely empty. The two primary types of chemical bonds are: ionic and covalent. Ionic Bonds With the exception of the noble gases, a neutral atom of almost any element is not particularly stable. It tends to gain or lose electrons in order to reach a stable configuration. When it does so, it becomes an ion. If it loses electrons it becomes a positively charged cation; if it gains electrons it becomes a negatively charged anion. Ions of equal but opposite charge are attracted to one another and form ionic bonds. A common example is sodium chloride (NaCl), table salt. Sodium, with an electron configuration of 1s2 2s2 2p6 3s1, would be more stable if it lost its 3s electron and became a cation; its charge would be +1. Chlorine, with a configuration of 1s2 2p6 3s2 3p5, would be more stable if it gained an electron to fill its 3s shell to its capacity of 6; it would be an anion with a charge of -1. Sodium's +1 charge and chlorine's -1 charge are equal and opposite, thus they attract one another. Sodium's one 3s electron joins chlorine's 3p shell to fill it. Both atoms have achieved a stable configuration and a strong bond. Atoms that become anions (i.e. gain electrons) are said to be highly electronegative; those that become cations (i.e. lose electrons) are said to have low electronegativity.

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The simplest ionic bond is one between two monoatomic ions. However, most ions are polyatomic, i.e. made up of multiple types of atoms. Examples are ammonium, bicarbonate, sulfate, and phosphate. The following table lists many common polyatomic ions along with their formula and charge.

Common Polyatomic Ions name

formula

charge

acetate

C2H3O2

-1

ammonium

NH4

+1

bicarbonate

HCO3

-1

bisulfate

HSO4

-1

carbonate

CO3

-2

chlorate

ClO3

-1

chlorite

ClO2

-1

chromate

CrO4

-2

cyanide

CN

-1

dichromate

Cr2O7

-1

hydroxide

OH

-1

mercury

Hg2

+2

nitrate

NO3

-1

nitrite

NO2

-1

permanganate

MnO4

-1

phosphate

PO4

-3

sulfate

SO4

-2

sulfite

SO3

-2

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An ionic bond can form when the ionic charges are equal. An example would be the cleaning compound trisodium phosphate, Na3PO4. Each of the three sodium ions has a charge of +1; the phosphate ion has a charge of -3. So, the charges are equal and the bond is formed. Compounds formed by ionic bonds are generally named with the cation first, followed by the anion. If the anion is a single element, the syllable -ide is added, as in sodium chloride. Most ionic compounds have very high melting points and are good conductors of electricity when melted or dissolved in water. The ionic attraction is between masses of atoms and there are not discrete molecules; rather, they tend to form a crystalline structure.

Crystalline structure of NaCl: each Na+ ion is surrounded by six Cl- ions, and each Cl- is surrounded by six Na+ ions.

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Covalent Bonds Compounds formed by covalent bonds do form discrete molecules. They are generally nonmetallic elements, and involve the sharing of electrons rather than transfer as in ionic bonds. They tend to have low melting points and do not conduct electricity. The simplest covalent bond occurs between two hydrogen atoms to form the H 2 molecule. Each hydrogen atom has a single electron, with room in its orbital for one more. When two hydrogen atoms are close enough together, their orbitals overlap and form a single molecular orbital. Both electrons now orbit both nuclei, and a stable covalent molecule has been formed. This bond can be graphically represented as follows:

H

H

HH

HH

This method of representing covalent bonds was developed by American chemist G.N. Lewis, and the electron dot diagrams are called Lewis structures. Lewis structures rely upon the octet rule, which states that atoms tend to be most stable when their valence (outermost) shell contains eight (or two for s-block elements) electrons. The dot diagrams show the number of electrons in the atom's valence shell. That number subtracted from eight if it is more than four, or the number itself if it is less than four gives the element's covalency, the number of electrons available for bonding.

Another simple covalent compound, water, illustrates this. Oxygen's Lewis structure is as follows:

O

Its covalency is two. Hydrogen's covalency is one, therefore two hydrogen atoms can bond

with one oxygen atom to form water.

HOH

The shared electron pair can also be noted as a

dash. Thus, a water molecule could also be noted as H--O--H. It is possible for atoms to share more than one pair of electrons and to form multiple covalent bonds. The oxygen molecule, O2, is an example of a double bond. As we have seen, oxygen's covalency is two. The two atoms share two pairs of electrons. The double bond can be represented by two dashes: O=O. Nitrogen, N 2, is an example of a triple bond. Its covalency is three, so the two atoms share three pairs of electrons. The triple bond can also be represented by three dashes: N≡N.

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O O

N

oxygen molecule O2

nitrogen atom

N

N

nitrogen molecule N2

Covalent compounds are named according to their chemical formula. The first word of the name is the central atom of the molecule, usually the least abundant element in the formula. The second word of the name is the other atom in the compound, usually with the suffix -ide added. If there are multiples of the same atom, a prefix such as di- (2), tri- (3), tetra- (4), penta- (5), or hexa(6) is added. Thus, CO2 is carbon dioxide. In some covalent compounds the electrons are not shared equally. The bond in this case is called a polar covalent bond and the compound is a polar compound. Polarity of bonds is a spectrum, not a fixed category. Ionic bonds are the most highly polar, and the covalent bonds between identical atoms (as in O2) are nonpolar. The degree of polarity is determined by the electronegativity of the elements that are bonded. Generally, elements to the right of the periodic table are more electronegative than those to the left. If the difference between the electronegativities of the two elements is greater than 1.7, the bond is considered ionic rather than covalent. Research into electronegativity and polarity of bonds earned Linus Pauling the first of his two Nobel Prizes in 1954. CHEMICAL FORMULAS A chemical formula is simply an ingredient list for a compound. It tells the types of atoms that are present and the quantity of each, in the form of a ratio. Thus, common salt, NaCl, has one sodium atom for every chlorine atom. Water, H2O, has two hydrogen atoms for each oxygen atom. The formula weight or molecular weight is the sum of the atomic masses of the atoms in the formula. So, the formula weight for NaCl is 58.44277; the sum of sodium's atomic mass, 28.98977, and chlorine's atomic mass, 35.453. The molecular weight of water is (1.0079)2 (twice hydrogen's atomic mass) plus 15.9994 (oxygen's atomic mass), or 18.0152. CHEMICAL EQUATIONS Chemical processes are called reactions and are expressed as equations. The equations describe the conditions of the reaction and how the atoms and molecules are recombined.

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The Mole Quantities of atoms and molecules are expressed in terms of moles. One mole is defined as the number of particles of a substance equal to the number of atoms in twelve grams of C-12. By various experimental methods that number has been determined to be 602,000,000,000,000,000,000,000, or 6.02 x 1023. This is called Avogadro's number. One mole of any substance weighs a number of grams equal to its formula weight. Thus, one mole of NaCl weighs 58.44277 grams. The mole is used in calculating quantities for chemical equations and concentrations of solutions. Calculations Using

the

mole

concept

along

with

chemical

formulas and formula weights, one can calculate the quantities needed of each reactant to yield a desired

product.

All

of

these

calculations

are

collectively called stoichiometry. For example, the compound FeS (iron II sulfide, used in treating anemia) is made by the following reaction: Fe + S → FeS. The formula tells us that one atom of iron is needed to react with each atom of sulfur. But since individual atoms are too small to be measured, we can say that one mole of iron is needed for each mole of sulfur. From their atomic masses we then know that we need 55.85 grams of iron and 32.07 grams of sulfur, and the resulting one mole of the compound would weigh 87.92 grams. Of course, if we need more than 88 grams of FeS, we can multiply the mole weights. If we need 500 grams of FeS, we divide the desired quantity (500g) by the mole weight (87.92) to find the number of moles of FeS--5.687. From the formula we know that one mole of each element combines to make one mole of compound, thus to obtain 5.687 moles of FeS we need that many moles of each element. Multiplying the mole weight of each element by 5.687 tells us that we need 317.62 grams of iron and 182.38 grams of sulfur. We can also calculate the mass percentage of each element. 317.62 grams of iron make up 63.5% of the mass of the 500 grams of FeS (317.62 divided by 500).

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Writing and Balancing Equations Chemical equations describe chemical processes or reactions. A variety of symbols are used. Often the physical state of each compound (solid, liquid, or gas) is indicated, by the letter (s), (l), or (g) beside the name of the compound. If heat is required to bring about the reaction, ∆ is placed above the reaction arrow to indicate that. An example is the decomposition of solid calcium carbonate (CaCO3) into solid calcium oxide (CaO) and gaseous carbon dioxide (CO 2) by heating. The equation is expressed as follows: CaCO3 (s) →∆ CaO (s) + CO2 (g). Just like mathematical equations, chemical equations must be balanced; i.e. the same number of atoms of each element must be present on each side of the equation. This is required by the Law of Conservation of Mass. The above example is balanced. Sometimes, however, we need to multiply quantities to achieve balance. For example, to combine sodium and oxygen to yield disodium oxide, you might write the equation as Na + O2 → Na2O. However, the equation isn't balanced. There is one sodium atom and two oxygen atoms on the left side of the equation, and two sodium atoms and one oxygen atom on the right side. In order to balance the equation, note the ratio in the product formula--two sodium atoms to one oxygen atom. Oxygen is a diatomic molecule--it always occurs as O2, so we can't decrease that quantity, but we can multiply the sodium. Four sodium atoms to two oxygen atoms meets the ratio, and would yield two molecules of disodium oxide. So, the balanced equation would be written as 4Na + O 2 → 2Na2O. SOLUTIONS AND SOLUBILITY A solution is a homogeneous mixture of two or more substances which do not tend to separate from the solution unless made to do so by physical means. No chemical bonds are broken when solutions are separated. An example would be separating a solution of salt and water by evaporating or distilling away the water to leave behind pure salt. A homogeneous mixture is one in which all parts of the mixture have the same composition.

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In any solution with two components, one is the solvent and the other is the solute. In the case of a solid and a liquid, such as salt in water, the liquid is the solvent and the solid is the solute. Usually the solvent is the most abundant component, though exceptions occur. In almost any solution with water, water is considered the solvent even if it is not the most abundant component. An example would be a 70% solution of rubbing alcohol in water.

The process of mixing a solvent and solute is called dissolving. The solute is dissolved and spread throughout the volume of the solvent. The ability of a solute to dissolve in a solvent is called solubility. Solubility is expressed as the number of grams of solute that can be dissolved in 100 ml of solvent. Some solutes have infinite solubility, such as ethyl or isopropyl alcohol in water. There is no limit to the quantity of these alcohols which can be dissolved in 100 ml of water. Sodium chloride has a solubility of 36 grams in 100 ml of water. Various factors affect solubility. The most important is the degree of polarity of the compound. The more polar the bond, the more soluble it is. Temperature also affects solubility. Sugar (sucrose) dissolves more readily in hot water than in cold water.

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Molarity The concentration of a solution tells how much solute is present in a given quantity of solution or solvent. Concentration is usually expressed as molarity. Molarity is defined as the number of moles of solute dissolved in one liter of solution; it is written as the capital letter M after the number. For example, to make one liter of a 1M solution of NaCl in water, weigh out one mole of NaCl (58.5 g.) and place it in a volumetric flask. Then fill the flask with water up to the 1-liter mark. To make a smaller quantity of solution, simply calculate the proportions. To make only 250 ml (1/4 liter) of the 1M NaCl, you need 1/4 mole of solute, or 14.6 g. Molarity calculations can also be used to change the concentration of a solution. Many prepared chemical solutions are sold in concentrated form. For example, HCl (hydrochloric acid) might be sold as 1 liter of 1M solution. If you need 250 ml of a 0.1M solution, first calculate the number of moles of HCl needed for that concentration. For one liter of 0.1M solution you would need 0.1 moles. For 250 ml you would need 0.025 moles. You know that you have one mole of HCl in 1000 ml, so multiply 1000 by 0.025 to get the quantity of solution that gives you 0.025 moles of HCl: 25 ml. Place 25 ml of the 1M HCl into a volumetric flask, then add water up to the 250 ml mark. The formula M1V1=M2V2 makes calculations simpler, where M1, M2 and V1, V2 are the molarities and volumes respectively. A closely related unit of concentration is molality. Molality is defined as the number of moles of solute dissolved in one kilogram of solvent. It is written as a lowercase m after the number.

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TYPES OF REACTIONS There are four primary types of chemical reactions, based on the types of reactants and products: combination, decomposition, displacement, and double displacement. A combination reaction involves the combining of two reactants to form a single new product. It can be generically written as A + B → A-B. An example is C + O2 → CO2. A decomposition reaction occurs when a compound breaks down into two or more elements or new compounds. Its generic form is A-B → A + B. An example is hydrogen peroxide decomposing to water and free oxygen: 2H2O2→ 2H2O + O2. In a displacement reaction, one substance reacts with a compound and replaces one of its elements. Generically it is written as A + B-C → A-B + C. An example is 2NaBr(s) + Cl2(g) → 2NaCl(s) + Br2(l). In a double displacement reaction, two compounds react with each other, and one or more atoms or groups of atoms from each exchanges with the other. The generic form of a double displacement reaction is A-B + C-D → A-D + B-C. An example is a medical test for the presence of calcium in urine. The test reagent ammonium oxalate, (NH4)2C2O4, is added. If calcium ions are present, the solid calcium oxalate will form. The reaction is as follows: CaCl2 + (NH4)2C2O4→ CaC2O4 + 2NH4Cl.

Redox Reactions Another way to classify types of reactions is to look at electron transfer during the reaction. This is determined by examining the oxidation numbers of the elements involved. These types of reactions are called oxidation-reduction or redox reactions. They are required in many biochemical processes such as converting food into usable energy for body functions. The oxidation number is a charge assigned to an atom by assuming that the bonded electrons are nearest to the more electronegative atom of the compound. In ionic compounds, the oxidation number is the same as the charge of the ion. In covalent compounds, the more electronegative atom will have a negative charge because it "pulls" the electron pair toward itself even while sharing them with the less electronegative atom. The number will be the number of shared electrons. The less electronegative atom will have a positive charge with a value equal to the

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number of electrons it is sharing. The oxidation numbers in covalent compounds can be easily determined from the Lewis dot structure of the molecule. In water, for example, the dot structure

is

HOH

. Each hydrogen atom is sharing one electron with oxygen, and oxygen is more

electronegative. Therefore hydrogen's oxidation number is +1. Oxygen is sharing two electrons, so its oxidation number is -2. The oxidation number of any given element can be different in different compounds. When the oxidation numbers of the reactants change as a result of the reaction, the reaction is a redox reaction. For example, when carbon and oxygen combine to form carbon dioxide, each element starts with an oxidation number of zero. In the product, CO2, carbon has an oxidation number of +4 because it has lost four electrons to the two oxygen atoms. Each oxygen atom gains two electrons, so its oxidation number is -2. Therefore the reaction C + O2 → CO2 is a redox reaction. Redox reactions are usually reversible, so they are often written with a double arrow ↔. In a redox reaction, one compound is oxidized and another is reduced. When an atom gains electrons its oxidation number becomes negative; it is called an oxidizing agent and it is said to be reduced itself. When it loses electrons, its oxidation number becomes positive; it is a reducing agent and it is said to be oxidized. Redox reactions are often represented as two half-reactions, showing the reduction and the oxidation separately.

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Exercises 1. Write the names for these ionic compounds: Compound formula

Compound name

NaHCO3 CaCO3 NH4NO3 Cu2SO4 AgNO3 KI

2. Write the Lewis electron dot structures for these compounds: Compound name

Compound formula

ammonia

NH3

hydrochloric acid

HCl

carbon dioxide

CO2

carbon tetrachloride

CCl4

Lewis dot structure

3. Name these covalent compounds: Compound formula

Compound name

CF4 PCl3 SF6 H2S

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4. Write the formulas for these compounds: Compound name

Compound formula

hydrogen bromide silicon dioxide oxygen dichloride silicon tetrahydride

5. Calculate the formula weight for these compounds, then the weight of one mole of each: Compound formula

Formula weight

Weight of 1 mole

Ba(NO3)2 FeSO4 H2O2 CH3NH2

6. What is the mass in grams of each of the following elements or compounds? Quantity of substance

Mass in grams

2.00 moles of N2 15.8 moles of SO2 1 mole of LiCl 1.86 moles of Au

7. What is the mass percent of each element in the compound NH4NO2?

8. Write a balanced equation to describe this reaction: sulfur dioxide gas reacts with oxygen gas to produce sulfur trioxide.

9. Balance this equation: H2SO4 + KOH → H2O + K2SO4

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10. Calculate the molarity of each of these solutions: Solution

Molarity

3.0 liters of solution containing 120 g of acetic acid (C2H3O2H) 20 g NaOH in 500 ml of solution 1 mole of solute in 500 ml of solution

11. Calculate the mass in grams of solute needed for these solutions: Solution

Grams of solute

1 L of 0.5M NaOH 100 ml of1M KCl 1 L of 0.3M NaCl 2 L of 0.1M glucose (C6H12O6)

12. What type of reaction is each of the following: reaction

type

2KClO3→ 3O2 + 2KCl ZnO + H2 → Zn + H2 CaO + CO2 → CaCO3 H2SO4 + 2KOH → 2H2O + K2SO4

13. What is the oxidation number for the colored element in each of these compounds? compound

oxidation number

CH4 P4 CaCO3 MnCl2

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14. Determine whether each of the following reactions is a redox reaction. reaction

redox

not redox

4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) Cl2(g) + H2O(l) → HCl(aq) + HClO(aq) 3As2O3 + 7H2O + 4NO3- + 4 H+↔ 6H3AsO4 + 4NO

15. In each of the following reactions, determine which reactant is the oxidizing agent and which is the reducing agent. reaction

oxidizing agent

reducing agent

2Mg + O2→ 2MgO Mg + Cl2→ MgCl2 Sn + F2→ SnF2 ZnO + H2→ Zn + H2O

Questions to Think About 1. How are ionic and covalent bonds similar? How are they different?

2. Why do you think the mole is used instead of counting individual atoms or molecules?

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