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CHAPTER 8 CHEMISTRY REVIEW

CHAPTER SUMMARY

This chapter is a general outline and review of the important chemistryconcepts that are tested by many nursing school entrance exams

Some of these key concepts are atomic structure, the periodic table, chemicalbonds, chemical equations, stoichiometry, energy and states of matter, reactionrates, equilibrium, acids, bases, oxidation-reduction, nuclear chemistry, andorganic compounds

B How to Use This Chapter

This chapter is presented in outline format as a systematic presentation of

important chemistry topics to help you review for your exam This does notconstitute a comprehensive chemistry review—use it as an aid to help you recallconcepts you have studied and to identify areas in which you need more study

At the end of this chapter, you will find a list of references and resources for amore complete review

Read each topic and answer the questions that follow After answering thesample test questions, you can pinpoint where you want to concentrate yourefforts If a question poses particular difficulty for you, study more problems ofthis type The more you hone your problem-solving skills, understand basic

STUDY TIPS FOR CHEMISTRY

• Review the topics covered in this chapter carefully Keep a copy ofone or more of the suggested resource books handy for more

extensive review

• Don’t try to review all topics in one or two study sessions Tackle acouple of topics at a time Focus more in-depth study on the itemswithin a topic about which you feel least confident first

• Complete each group of practice questions after you study eachtopic, and check your answers If you experience particular difficultywith one type of question, choose similar questions from the otherresources listed to practice some more

• Review all the answer choices carefully before making your

selection The wrong answers often give you hints at the correct oneand help you confirm that you really do know the correct answer.Remember that recognition is not necessarily understanding

• When checking your answers to practice questions with the answerkey, be sure you understand why the identified choice is the correctone Practice writing out your reasoning for choosing a particularanswer and checking it against the reasoning given in the answerkey

• Practice pronouncing chemical terminology aloud If you can

pronounce a term with ease, you are more likely to remember theterm and its meaning when reading it

• Review carefully the visual aspects of chemistry, such as the use ofsymbols, arrows, and sub- and superscripts If you know the

circumstances under which particular symbols are used, you willhave immediate clues to right and wrong answers

• Focus on developing problem-solving skills Almost all chemicalproblems require the analysis, sorting, and understanding of details

II Main Topics

Greek atomos, meaning indivisible).

• Atoms of one element are identical in size, mass, and chemical properties

• Atoms of different elements have different sizes, masses, and chemicalproperties

• Chemical compounds are made up of atoms of different elements in a ratiothat is an integer (a whole number) or a simple fraction

Atomic number is the number of protons in the atom and is specific for each

element The atomic number is indicated by the number to the lower left ofthe element symbol: 11Na

5 Isotopes

Isotopes are atoms of the same element that have the same number of protons

(same atomic number) but different number of neutrons (different mass

number) Isotopes have identical chemical properties (same reactivity) butdifferent physical properties (for example, some decay while others are

In nature, atoms of one element may be chemically bonded to other atoms

of the same element For example, hydrogen and oxygen are always diatomic,which means that they naturally exist as H2 and O2, respectively Elementalsulfur exists as S8 Many elements, like sodium, exist as single atoms in theirelemental form

7 Properties of Atoms

Law of conservation of mass: In a chemical reaction, matter cannot be

created or destroyed—i.e., the mass of the reagents always equals the mass ofthe products Likewise, the number of each type of atom will be equal oneach side of the reaction

Law of constant (definite) proportion: A chemical compound will

always have the same proportion of elements by mass—e.g., water (H2O) willalways be 8/9 oxygen and 1/9 hydrogen by mass

Law of multiple proportions: If two elements form more than one

compound between them, then the ratios of the masses of the second elementwhich combine with a fixed mass of the first element will be ratios of smallwhole numbers For example, 16 g of oxygen will react with 14 g of nitrogen

to form NO and 28 g of nitrogen to form N2O (1:2 ratio)

Questions

of valence electrons Elements in the same group share similar chemicalproperties

c Metals

A metal is an element that is a good conductor of heat and electricity in

addition to being shiny (reflecting light), malleable (easily bent), and ductile

d Nonmetals

A nonmetal is an element with poor conducting properties They are

electronegative and accept electrons in their valence shell They are found inthe upper right-hand corner of the periodic table (group VIIA, parts of IVA-VIA)

e Metalloids

A metalloid is an element with properties that are intermediate between those

of metals and nonmetals, such as semiconductivity They are found betweenmetals and nonmetals in the periodic table

3 Electronic Structure of Atoms

a Bohr Atom

Niels Bohr’s planetary model of the hydrogen atom, in which a nucleus issurrounded by orbits of electrons, resembles the solar system Electrons could

be excited by quanta of energy and move to an outer orbit (excited level).

They could also emit radiation when falling to their original orbit (groundstate)

the electron n is always a positive integer (n = 1, 2, 3, …) For a givenprinciple quantum number (n), there are n – 1 possible orbital quantumnumbers (0, 1, 2, … , n – 1) The principle quantum number defines theenergy level of an electron There are a maximum of n2 orbitals and 2n2electrons in an energy level

The outer shell is the last energy level in which loosely held electrons are

contained These are the electrons that engage in bonding and are thereforecharacteristic of the element

You Should Review

• periodic table: structure; specific names of the different groups (group IA:alkali metal, group IIA: alkaline earth, group VIIA: halogens, etc.); thelocation of metals, nonmetals, and metalloids

A cation results when an atom loses one or more electrons, becoming

positively charged Most cations are metallic and have the same name as themetallic element For example, lithium ion (Li+) has one electron less thanlithium atom (Li), having acquired the noble gas electron structure of helium,

The VSEPR model is based on electrostatic repulsion between electron pair

orbitals By pushing each other as far as possible, electron pairs dictate whichgeometry, or shape, a molecule will adopt Molecules should be written asLewis structures (see the preceding electron–dot notation)

d Electronegativity and Dipoles

Electronegativity is the ability of an atom in a bond to attract the electron

density more than the other atom(s) in the bond Electronegativity increasesfrom left to right and from bottom to top in the periodic table Thus, fluorine(F) is the most electronegative element of the periodic table, with the

maximum value of 4.0 in the Pauling scale of electronegativity The Paulingscale is a range of electronegativity values based on fluorine having the

highest value at 4.0 These values have no units Metals are electropositive,with a minimum electronegativity value of 0.8 on the Pauling scale for mostalkali metals

A dipole results in a covalent bond between two atoms of different

electronegativity Partial positive (+δ) and negative (–δ) charges develop at

Electronegative atoms (such as F, N, or O) covalently bonded to H atomsare considered hydrogen bond donors Electronegative atoms with free lonepairs of electrons in their Lewis structures act as hydrogen bond acceptors

Acetone ((CH3)2CO) is a hydrogen bond acceptor Ammonia (NH3) is ahydrogen bond donor and acceptor Like ammonia, water is both a hydrogenbond donor and acceptor

6 Polyatomic Ions

Polyatomic ions are groups of two or more covalently bonded atoms thatpossess a positive or negative charge They form ionic compounds in thesame way as single-atom ions Polyatomic ions can be as simple as hydroxide(OH–) Other common examples are ammonium , phosphate

carbonate nitrate and sulfate

b AM+, Ag+, CO32–, NO3–, Cal+, Fl–

c NH4–, Ag+, HCO3–, NO3–, Cal+, Fl–

d NH4+, Ag+, HCO3–, NO3–, Ca2+, F–

Answers

21 b A covalent bond exists between H and O in the H2O molecule Since thebond is formed between two elements with different electronegativities, it

two electrons

26 a The electron configuration of element 19 is 1s22s22p 63s 23p 64s1 Toachieve the outer octet, it must give away one electron, thus gaining a

molecular formula It is the same as the molar mass (in grams) without theunit

2 Moles

A mole of a particular substance is defined as the amount that contains the

number of atoms in exactly 12 g of carbon-12 Experiments have establishedthat number to be 6.02214199 1023 particles per mole (Avogadro’s

32 The formula of carbon dioxide is CO2 Its molecular weight is 44 amu Asample of 11 grams of CO2 contains

3 C = 3 12 = 36 amu

9 O = 9 16 = 144 amu

Then add them up to get the molecular weight, which is 234 amu

32 c

33 b 1 mole of Ca3(PO4) 2 = 310 g; 0.200 mol 3 310 g/mol = 62 g

34 c 5O2 = 5 mol 3 32 g/mol = 160 g

35 d One mole of CaCl2 would be needed to get 2 mol NaCl Since 0.5 mol ofNaCl, or 25% of 2 moles, is to be formed, 0.25 mol CaCl2 (25% of 1 mole)

• Gas molecules collide with each other; they do not attract or repel eachother

c Gay-Lussac’s Law (at constant volume)

The pressure of any sample of gas maintained at constant volume increaseswith the temperature Temperature needs to be converted from Celsius (°C) toKelvin (K) when using this law (K = °C + 273)

d Avogadro’s Law (at constant T and P)

The volume of gas increases with the number of moles of gas present at

constant temperature and pressure

Standard temperature and pressure (STP) is achieved at 273 K and 1 atm(760 torr) when one mole (or 6.023 1023 particles) of any gas occupies a

An ideal gas is a gas whose pressure, volume, and temperature obey the

relation, PV = nRT (a combination of Boyle’s, Charles’s, and Avogadro’slaws), with R being the gas constant Temperature needs to be converted fromCelsius (°C) to Kelvin (K) when using this law (K = °C + 273) The samerelationship can also be expressed as:

Solubility of most solids and liquids increases with increasing temperaturewhile decreasing for gases dissolved in liquids (gas molecules tend to

escape)

3 Concentration of Solutions

Percent concentration expresses the concentration as a ratio of the weight (orthe volume) of the solute over the weight (or the volume) of the solution Thisratio is then multiplied by 100

Weight/volume% = grams of solute/100 ml of solvent

Volume/volume% = volume of solute/100 ml volume of final solution

Weight/weight% = grams of solute/100 g of solution

4 Molarity

Molarity (M) expresses the number of moles of solute per liter of solution A

0.1 M NaOH aqueous (dissolved in water) solution has 0.1 mol of solute(NaOH) per 1 liter of water

5 Dilution

M iVi = M f V f (i = initial; f = final) establishes the equivalence between the

initial and final concentrations In dilution, equivalence must be achievedbetween the initial and final concentrations

Since M (mol/L) V (L) gives units of moles, this equation states that theamount of a substance must be constant before and after a dilution occurs,i.e., if 1 L of an aqueous solution containing 0.1 mol (5.8 g) of NaCl is diluted

by adding an additional liter of water, there will still be 0.1 mol (5.8 g) ofNaCl in the solution

hydrogen bonds It is found in large amounts in cells and blood Water is anexcellent solvent and has a high boiling point, high surface tension, high heat

Specific heat of vaporization is the heat required to evaporate 1 gram of a

liquid Water’s large heat of vaporization (540 calories/gram) requires large

amounts of heat in order to vaporize it During perspiration, water evaporatesfrom the skin, and large amounts of heat are lost.

53 c 250 ml 0.2 mol NaOH/1,000 ml = 0.05 mol; 0.05 mol 40 g/mol = 2.00 g

54 a 250 ml 0.3 mol NaCl/1,000 ml = 0.0750 mol

55 a The other properties listed are due to the polar nature of water.

56 d.

The equilibrium constant, K, for a reaction describes the concentrations of

reactants and products for a chemical reaction at equilibrium K is often

dependent on temperature For a balanced chemical equation, ,the equilibrium constant is written as:

where [A], [B], [C], and [D] are concentrations of reactants and products and

w, x, y, and z are the coefficients used to balance the chemical equation If one

of the reactants or products is a solid or pure liquid (not a solution), it is notincluded in the equilibrium expression

3 Activation Energy

be transformed into products (i.e., to overcome the energy barrier betweenreactants and products) The higher the activation energy, the slower the

b Particle Size

Smaller particles react faster, as they collide often at any given temperatureand concentration

c Concentration

A high concentration of reacting particles increases the rate of chemical

d Catalysis

Catalysts speed the reaction rate by lowering the activation energy of thereaction They are not consumed in the reaction

6 Reversible Reactions

A double arrow designates reversible (two-way) chemical reactions Ifarrows differ in length, the longer arrow indicates the major (faster) direction

You Should Review

• Le Chatelier’s principle and the different stresses that can be placed onchemical processes

completely dissociate in water Acids have a sour taste.

dissolved in water, strong bases such as NaOH dissociate to release hydroxideions and sodium cations Bases have a bitter taste and feel slippery like soap

2 Reactions of Acids

Common reactions include:

• metal + acid salt + hydrogen Zn + 2HCl ZnCl2 + H2

• base + acid salt + water NaOH + HNO3 NaNO3 + H2O

• metal oxide + acid salt + water CaO + 2HNO3 Ca(NO3)2 + H2O

• metal carbonate + acid salt + carbonic acid (unstable) NaHCO3 + HCl NaCl + H2CO3 (H2CO3 H2O + CO2)

3 Autoionization of Water

[H3O+] = [OH–]

The ion product of water is Kw: Kw = [H3O+] [OH–] = 1 10–14 Thus,

in pure water: [H3O+] = [OH–] = 1 10–7 moles/liter Note that H3O+

(hydronium ion) is often abbreviated as H+ (hydrogen ion)

4 pH

pH = – log [H + ] The pH measures the negative logarithm (for presentation

of very small numbers in a large scale) of the hydrogen ion concentration (inmoles/liter) The pH scale runs from 0 to 14, with acids in the lower end ofthe scale (less than pH 7) and bases at the higher end (greater than pH 7)