Chapter 3 Matter—Properties and Changes (Ruiki Basilio, Angeles Phan, Albert Alavarez and Carlos Quintanilla)

3.1 Properties of Matter

A substance is a form of matter with a uniform and unchanging composition.

Physical properties can be observed without altering a substance’s composition. Chemical properties describe a substance’s ability to combine with or change into one or more new substances.

Both physical and chemical properties are affected by external conditions such as temperature and pressure.

The three common states of matter ate solid, liquid, and gas.

Review Questions:

List three examples of substances. Explain why each is a substance.

Classify each of the following as a physical property or a chemical property.a. aluminum has a silvery colorb. gold has a density of 19g/cm3

c. sodium ignites when dropped in waterd. water boils are 100°Ce. silver tarnishesf. mercury is a liquid at room temperature

3.2 Changes in Matter

A physical change alters the physical properties of a substance without changing its composition.

A chemical change, also known as a chemical reaction, involves a change in a substance’s composition.

In a chemical reaction, reactants form products.

The law of conservation of mass states that mass sis neither created nor destroyed during a chemical reaction; it is conserved.

Review Questions:

Describe the difference between a chemical change and a physical change.

Classify each of the following as a physical change or a chemical change.a. Breaking a pencil in twob. Water freezing and forming icec. Frying an eggd. Burning woode. Leaves turning color in the fall

3.3 Mixtures of Matter

A mixture is a physical blend of two or more pure substances in any proportion.

Solutions are homogenous mixtures.

Mixtures can be separated by physical means. Common separation techniques include filtration, distillation, crystallization, and chromatography.

Review Questions:

Describe how a homogeneous mixture differs from a heterogeneous mixture.

Describe a method that could be used to separate each of the following mixtures.a. Iron filings and sandb. Sand and saltc. The components of inkd. Helium and oxygen gases

3.4 Elements and Compounds

Elements are substances that cannot be broken down into simpler substances by chemical or physical means.

The elements are organized in the periodic table of elements.

A compound is a chemical combination of two or more elements. Properties of compounds differ from the properties of their component elements.

The law of definite proportion states that a compound is always composed of the same element in the same proportions.

The law of multiple proportions states that if elements form more than one compound, those compounds will have compositions that are small, whole-number multiples of each other.

Review Questions:

Is it possible to distinguish between an element and a compound? Explain.

Name the elements contained in the following compounds.a. Sodium chloride (NaCl)b. Ammonia (NH3)c. Ethanol (C2H6O)d. Bromine (Br2)

Chapter 4 The Structure of the Atom (Gonzalez Ileana & Barrios Cindy)

4.1 Early Theories of Matter

The Greek philosopher Democritus was the first person to propose the existence of atoms.

In 1808, Dalton proposed his atomic theory, which was based on numerous scientific experiments.

All matter is composed of atoms. An atom is the smallest particle of an element that maintains the properties of that element. Atoms of one element are different from atoms of other elements.

4.2 Subatomic Particles and the Nuclear Atom

Atoms are composed or negatively charged electrons, neutral neutrons, and positively charged protons. Electrons have a 1 – charge, protons have a 1+ charge, and neutrons have no charge. Both protons and neutrons have masses approximately 1840 times that of an electron.

The nucleus of an atom contains all of its positive charge and nearly all of its mass.

The nucleus occupies an extremely small volume of space at the center of an atom. Most of an atom consists of empty space surrounding the nucleus through which the electrons move.

Atomic Number = number of protons; number of protons = number of electrons Mass number - Atomic number = number of neutrons (mass number is also known as atomic

mass)

Review Questions:

1. An atom is composed mainly of:A) ProtonsB) ElectronsC) Empty SpaceD) Water

2. True or FalseThe smallest particle of an element that retains the properties of the element is called an atom

3. Electrons are located in A) The nucleusB) The space surrounding the nucleusC) The space surrounding the atomD) In the nucleolus

4. The negatively charged particles are called?A) NeutronsB) ProtonsC) ElectronsD) Isotopes

Answer Key: C, True, B , C

4.3 How Atoms Differ

The number of protons in an atom uniquely identifies an atom. This number of protons is the atomic number of the atom.

Atoms have equal numbers of protons and electrons, and thus, no overall electrical charge.

An atom’s mass number (or atomic mass) is equal to its total number of protons and neutrons.

Atoms of the same element with different numbers of neutrons and different masses are called isotopes.

The atomic mass of an element is a weighted average of the masses of all the naturally occurring isotopes of that element.

4.4 Unstable Nuclei and Radioactive Decay

Chemical reactions involve changes in the electrons surrounding an atom. Nuclear reactions involve changes in the nucleus of an atom.

The neutron-to-proton ratio of an atom’s nucleus determines its stability. Unstable nuclei undergo radioactive decay, emitting radiation in the process.

Chapter 5 Electrons in Atoms (Aukura Williams, Valerie Zuniga, Jacqueline Rodriguez)

5.1 Light and Quantized Energy All waves can be described by their wavelength, frequency, amplitude. and speed. Light is an electromagnetic wave. In a vacuum all electromagnetic waves travel at the speed

of 3.00 x 108 m/s. All electromagnetic waves may be described as both waves and particles. Particles of light are

called photons. Energy is emitted and absorbed by matter in quanta. In contrast the continuous spectrum produced by white light, an element’s atomic emission

spectrum consists of a series of fine lines of individual colors.

5.2 Quantum Theory and the Atom According to the Bohr model of the atom, hydrogen’s atomic emission spectrum results from

electrons dropping from higher-energy atomic orbits to lower-energy atomic orbits. The de Broglie equation predicts that all moving particles have wave characteristics and

relates each particle’s wave length to its mass, its velocity, and Planck’s constant. The quantum mechanical model of the atom is based on the assumption that electrons are

waves. The Heisenberg uncertainty principle states that it is not possible to know precisely the

velocity and the position of a particle at the same time. Electrons occupy three-dimensional regions of space called atomic orbitals. There are four

types of orbitals, denoted by the letters s, p, d, and f.

Review Questions:

1. What is the complete electron configuration of a scandium atom?A. 1s22s22p63s23p64s23d1

B. 1s22s22p73s23p74s23d1

C. 1s22s22p53s23p54s23d1

D. 1s22s12p73s13p74s23d1

5.3 Electrons Configurations The arrangement of electrons in an atom is called the atom’s electron configuration. Electron

configurations are prescribed by three rules: the aufbau principle, the Pauli exclusion principle, and Hund’s rule.

Electrons related to the atom’s highest principle energy level are referred to as the valence electrons. Valence electrons determine the chemical properties of an element.

Electron configurations may be represented using orbital diagrams, electron configuration notation, and electron-dot structures.

Review Questions:

1. Which of the following orbitals has the highest energy?A. 4fB. 5pC. 6sD. 3d

Chapter 6 The Periodic Table and Periodic Law (Christopher Molina and Fortunato Rojas)

6.1 Development of the Modern Periodic Table

Period law states that when the elements are arranged by increasing atomic number, there is a periodic repetition of their chemical and physical properties.

Newland’s law of octaves, which was never accepted by fellow scientists, organized the elements by increasing atomic mass. Mendeleev’s periodic table, which also organized elements by increasing atomic mass, became the first widely accepted organization scheme for the elements. Moseley fixed the errors inherent in Mendeleev’s table by organizing the elements by increasing atomic number.

The periodic table organizes the elements into periodic (rows) and groups (columns) by increasing atomic number. Elements with similar properties are in the same group.

Elements are classified into metals, nonmetals, or metalloids. The stair-step line on the table separates metals from nonmetals. Metalloids border the stair-step line.

6.2 Classification of the Elements

Elements in the same group on the periodic table have similar chemical properties because they have the same valence electrons configuration.

The four blocks on the periodic table can be characterized as follows: S-block: filled or partially filled s orbitals P-blocks: filled or partially filled p orbitals D-block: filled outer most s orbital of energy level n and filled or partially filled d

orbitals of energy level n – 1. F-block: filled outermost s orbital, and filled or partially filled 4f and 5f orbitals.

For the group and elements, an atom’s group number equals its number of valence electron.

The energy level of an atom’s valence electrons equals its period number. The s2p6 electron configuration of the group 8A elements (noble gases) is exceptionally

stable.

6.3 Periodic Trends Atomic radii generally decrease as you move left to right across a period and increase as you

move down a group. Positive ions are smaller than the neutral atoms from which they form. Negative ions are

larger than the neutral atoms from which they form. Ionic radii of both positive and negative ions decrease as you move left to right across the

period. Ionic radii of both positive and negative ions increase as you move down a group. Ionization energy indicates how strongly an atom holds onto its electrons. After the valence

electrons have been removed from an atom, there is a tremendous in the ionization energy required to remove the next electron.

Ionization energies generally increase as you move left to right across a period and decrease as you move down a group.

The octet rule states that atoms gain, lose, or share electrons in order to acquire the stable electron configuration of a noble gas.

Electronegativity, which indicates the ability of atoms of an element to attract electrons in a chemical bond, plays a role in determining the type of bond formed between elements in a compound.

Electronegativity values range from 0.7 to 3.96, and generally increase as you move left to right across a period and decrease as you move down a group.

Review Questions:

1. What are groups and periods and how are they different? Answer: Groups are the columns in the periodic table and periods are the rows which is their difference.

2. Where are alkali metals and alkaline earth metals located in the periodic table?Answer: Alkali metals are located in the first group and alkaline earth metals are located in the second group.

3. What are valence electrons?Answer: These are the electrons in the outermost energy level.

4. Give the electron configuration of Al.Answer: 1s22s22p63s23P1

5. What does ionization energy and atomic radius have in common?Answer: Their similarity is that they both decrease left to right across the period.

6. What does it mean for an electron to have a high electronegativity?Answer: It means that it is able to attract any electrons to them when they bond with another atom.

Chapter 7 The Elements (Marlin Gramajo & Jazmine Gomez)

7.1 Properties of S-block Elements

The number and location of valence electrons determine an elements precision on the periodic table and its chemistry.

Properties within a group are not identical because members have different numbers of inner electrons.

Similarities between period two elements and period three elements in neighboring groups are called diagonal relationships.

The representative elements in groups 1A through 8A have only S and P electrons. Because hydrogen has a single electron, it can behave as a metal and loose an electron or

behave as a nonmetal and gain an electron, The alkali and earth metals in group 1A and 2A are the most reactant metals. Metals form mixtures called alloys whose composition can be adjusted to produce different

properties. Sodium and Potassium are the most abundant alkali metals. Many biological functions are

controlled by Sodium and Potassium ions. Calcium is essential for healthy teeth and bones. It is most often found as Calcium carbonate,

which can decompose to form lime-one of the most important industrial compounds. Magnesium is used in lightweight, yet strong alloys. Magnesium ions are essential for

metabolism, muscle function and photosynthesis.

Review Questions:

1. What are Group 1A elements called?A) Alkali metals*B) HalogensC) AlkalineD) Alkaline Earth metals

2. What are the most abundant alkali metals?A) franciumB) nitrogenC) sodium and potassium*D) calcium

7.2 Properties of P-block elements P-block elements include elements, metalloids, nonmetals, and inert gases. Aluminum is the most abundant metal in Earth’s crust. Much more energy is needed to

extract Aluminum from its ore to recycle aluminum. Because of carbon atom can join with up to four other carbon atoms, carbon forms millions of

organic compounds. Graphite and diamonds are allotropes of carbon. The most abundant element in Earth’s crust, silicon and oxygen, are usually found in silica ,

which can be melted and rapidly cooled to form glass.

Lead, which is still used in storage batteries, was used in pipes, paint, and gasoline until people realized the danger of lead poisoning.

Nitrogen combines with hydrogen in ammonia, which is used in cleaning products. Nitric acid, which is produced from ammonia, is used to make solid fertilizers, explosives, and dyes.

Phosphates in fertilizers and cleaning products can harm the environment. Sulfur dioxide reacts with water to form one of the acids in acid rain. Most sulfur dioxide is

used to make sulfuric acid. Halogens are extremely reactive nonmetals. Their compounds are used in toothpaste,

disinfectants, and bleaches. Many plastics contain chlorine. Silver bromide or iodide is used to coat photographic film.

The stable noble gases are used in lighter than air blimps, in neon lights, as a substitute for nitrogen in diving tanks, and as an inert atmosphere for welding.

Review Questions:

1. Which of the following groups is composed entirely of nonmetals?A. 1AB) 3AC) 5AD) 7A*

2. What are group 7A elements called?A) halogens*B) noblesC) alkaline metalsD. alkali metals

7.3 Properties of d-block and f-block elements The d-block transition metals and inner transition metals are more similar across the period

than are the s-block and p-block elements. The more unpaired electrons in d sublevel, the harder the transition metal and the higher is

melting and boiling points. Ions with partially filled sublevels often form compounds with color.

In ferromagnetic metals, ions are permanently aligned in the direction of a magnetic field. Many transition metals are strategic materials. Lanthanides are silvery metals with high melting points that are found mixed in nature and

are hard to separate. Actinides are radioactive elements.

Review Questions:

1. Which of the following is not a property of transition metals? A) malleabilityB) electrical conductivityC) high solubility in water*D) luster

2. Which physical properties of transition metals increase with increasing unpaired electrons in the d sublevel?

A) hardnessB) melting pointC) boiling pointD) all of the above*

Chapter 8 Ionic Compounds (Mary Salveron and Abby Vargas)

8.1 Forming Chemical Bonds

A chemical bond is the force that holds two atoms together.

Atoms that form ions gain or lose valence electrons to achieve the same electron arrangement as that of a noble gas, which is the stable configuration. This noble gas configuration involves a complete outer electron energy level, which usually consists of eight valence electrons.

A positive ion, or cation, forms when valence electrons are removed and a stable electron configuration is obtained.

A negative ion, or anion, forms when valence electrons are added to the outer energy level, giving the ion a stable electron configuration.

8.2 The Formation and Nature of Ionic Bonds

An ionic bond forms when anions and cations close to each other attract, forming a tightly packed geometric crystal lattice.

Lattice energy is needed to break the force of attraction between oppositely charged ions arranged in a crystal lattice.

The physical properties of ion solids, such as melting point, boiling point, hardness, and the ability to conduct electricity in the molten state and as aqueous solution, are related to the strength of the ionic bonds and the presence of ions.

An ionic compound is an electrolyte because it conducts an electric current when it is liquid in aqueous solution.

8.3 Names and formulas for ionic compounds (Jenny G. and Selvin T.)

Subscripts in an ionic compound indicate the ratio of cations and needed to form electrically neutral compounds. The formula unit represents the ratio of these ions in the crystal lattice.

If the element that forms the cation has more than one possible oxidation number. Roman numerals are used to indicate the oxidation number present for that element in the compound..

Ions formed from only one atom are monatomic ions. The charge on a monatomic ion is its oxidation number, or oxidation state.

Polyatomic ions are two or more atoms bonded together that act as a single unit with net charge. Many polyatomic ions are oxyanions, containing an atom, usually a nonmetal, and oxygen atoms.

In a chemical formula, polyatomic ions are placed inside parentheses when using a subscript.

Ionic compounds are named by the name of the cation followed by the name of the anion.

Review Questions:

Determine the correct formula for the ionic compound composed of the following pairs of ions:

1) sodium and nitrate (NO3-)

2) calcium and chlorate (ClO3-)

3) aluminum and hydroxide (OH-)

4) potassium and chromate (CrO42-)

8.4 Metallic bonds and properties of metals

Metallic bonds are formed when metal cations attract free valence electrons. A sea of electrons moves throughout the entire metallic crystal producing this attraction.

The electrons involved in metallic bonding are called delocalized electrons because they are free to move throughout the metal and are not attached to a particular atom

The electron sea model can explain melting point, boiling point, malleability, conductivity and ductility of metallic solids.

Metal alloys are formed when a metal is mixed with one or more other elements. The two common types of alloys are substitutional and interstitial

Review Questions:

1) What is a metallic bond?

Answer: The attraction of a metallic cation for delocalized electrons.

2) What is an alloy?

Answer: A mixture of elements that has metallic properties.

Chapter 9 Covalent Bonding (Pedro Aguilar and Jin Mei Situ)

9.1 The Covalent Bond

A covalent bond is formed when atoms share one or more pairs of electrons.

Molecules, formed when atoms share electrons, are more stable then their constituent atoms.

Sharing a single pair of electrons results in a single covalent bond. Two atoms sharing more than one pair of electrons results in a multiple bond.

A double covalent bond results when two pairs of electrons are shared between atoms. Sharing three pairs of electrons result in a triple covalent bond.

When an electron pair is shared by the direct overlap of parallel orbitals, a sigma bond results. The overlap of parallel orbitals forms a pi bond. Single bonds are sigma bonds multiple bonds involve both sigma and pi bonds.

Bond length depends on the sizes of the bonded atoms and the number of electron pairs they share. Bond dissociation energy is the energy needed to break a covalent bond. Bond length and bond dissociation energy are directly related.

Practice Problems: Draw the Lewis Structure for each of these molecules:1. PH3

2. CCl4

1 2

9.2 Naming molecules

Names of covalent molecular compounds include prefixes that tell the number of each atom present.

The molecules that produce hydrogen ions in solution are acids and are named accordingly.

Practice Problems: Name the following binary covalent compounds:1. As2O3

2. CO1* diarsenic trioxide2* carbon monoxide

9.3 Molecular Structure

The Lewis Structure is used to show the distribution of shared and lone pairs of electrons in a molecule.

Resonance occurs when more than one valid Lewis structure exists for the same molecule.

Exceptions to the octet rule occur when an odd number of valence electrons exists between the bonding atoms, not enough electrons are available for an octet, or more than eight electrons are shared.

Coordinate covalent bonding occurs when one atom of the bonding pairs supplies both shared electrons.

Practice Problems: Draw the correct Lewis Structure for the following molecules, which contain expanded octets:

1. PCl5

2. SF6

1 2

9.4 Molecular Shape (Anna Hu & Denisee Alonso)

The valence shell electron pair repulsion, or VSEPR, model can be used to predict the three-dimensional shape of a molecule. Electron pair repel each other and determine both the shape of and bond angles in a molecule.

Hybridization explains the observed shapes of molecule by the presence of equivalent hybrid orbitals.

Two orbitals from two sp hybrid orbitals, and the molecule is linear. Three orbitals, forming three sp2 hybrid orbitals, form a molecule that is trigonal planar. Four orbitals, forming sp3 hybrid orbitals, form a molecule that is tetrahederal.

9.5 Electronegativity and Polarity

Electronegativity is the tendency of an atom to attract electrons and is related to electron affinity. The electronegativity difference between two bonded atoms is used to determine the type of bond that most likely occurs.

Polar bonds occur when electrons are not shared equally, resulting in an unequal distribution of charge and the formation of a dipole.

The spatial arrangement of polar bonds in a molecule determines the overall polarity of a molecule.

Weak intermolecular forces, also called van der Waals forces, hold molecules together in the liquid and solid phases. These weak attractive forces determine properties. Molecular solids tend to be soft and have low melting and boiling points.

Covalent network solids result when each atom is covalently bonded to many other atoms in the solid. These solids are hard and have high melting points.

Practice Problems:

For each of the following, predict the molecular shape. (9.4)

a. COS b. CF2Cl2 Answers: a. linear b. tetrahedral

All of the following compounds have bent molecular shapes EXCEPT _________. (9.4)

a. BeH2 b. H2S c. H2O d. SeH2 Answer: a

Predict which of the following bonds is the most polar. (9.5)

a. C—O b. Si—O c. C—Cl d. C—Br Answer: a

Which of the following compounds is NOT polar? (9.5)

a. H2S b. CCl4 c. SiH3Cl d. AsH3 Answer: b

Chapter 10 Chemical Reactions (Daniel Campo and Darreyon Johnson)

10.1 Reactions and Equations

Some chemical reactions release energy in the form of heat and light, and some absorb energy.

changes in temperature, color, odor, and physical state are all types of evidence that indicate a chemical reaction has occurred.

Word and Skeleton equations provide important information about a chemical reaction, such as reactant and products involved in the reaction and their physical states.

A chemical equation gives the identities and relative amounts of the reactants and products that are involved in a Chemical Reactions. Chemical equations are balanced.

Balancing an equation involves adjusting the coefficient of the chemical formula in the skeleton equation until the number of each element is equal on both sides.

EXAMPLE

Step 1: Write the skeleton equation. Be sure to put the reactants on the left side of an arrow and the products on the right. Separate the substances with plus signs and indicate physical states.

NaOH (aq) + CaBr2 (aq) Ca(OH)2 (s) + NaBr (aq)

Step 2: Count the atoms of each element in the reactants.

1 Na, 1 O, 1 H, 1 Ca, 2 Br

Step 3: Count the atoms of each element in the products.

1 Na, 2 O, 2 H, 1 Ca, 1 Br

Step 4: Adjust Coefficients. Insert the coefficient 2 in front of NaOH to balance the hydroxide ions.

2NaOH + CaBr2 Ca(OH)2 + NaBr

Insert the coefficient 2 in front of NaBr to balance the Na and Br atoms.

2NaOH + CaBr2 Ca(OH) 2 + 2 NaBr

Step 5: Write the Coefficient in their lowest possible ratio. The ratio of the coefficient is 2:1:1:2.

Step 6: Check to make sure that the number of atoms of each element is equal on both sides of the equation.

Reactants: 2 Na, 2 OH, 1 Ca, 2 BrProducts: 2 Na, 2 OH, 1 Ca, 2 Br

Write the skeleton equation and balance the equation.

1.) hydrogen (g) + bromine (g) hydrogen bromide (g)

H2 (g) + Br2 (g) HBr (g)

Symbols:

+ --------Separates two or more reactants or products. -------Separates reactants from products.(s)-------identifies Solid(l) ------ Identifies Liquid State(g) ----- Identifies gaseous State(aq) ----Identifies water solution

10.2 Classifying Chemical Reactions (Kimberly Decastro & Heydy Hernandez)

Classifying chemical reactions makes them easier to understand, remember, and recognize.

Synthesis, combustion, decomposition, single-replacement, and double-replacement reactions are five classes of chemical reactions.

A synthesis reaction occurs when two substances react to yield a single product. The substances that react can be two elements, a compound and an element, or two compounds.

A combustion reaction occurs when a substance reacts with oxygen, producing heat and light. A decomposition reaction occurs when a single compound breaks down into two or more

elements or new compounds.

A single-replacement reaction occurs when the atoms of one element replace the atoms of another element in a compound.

In single-replacement reactions, a metal may replace hydrogen in water, a metal may replace another metal in a compound dissolved in water, and a nonmetal may replace another nonmetal in a compound.

Metals and halogens can be ordered according to their reactivities. These listings, which are called activity series, can be used to predict if single-replacement reactions will occur.

A double-replacement reaction involves the exchange of positive ions between two compounds.

Review Questions:

Write a chemical equation for the following decomposition reaction.

Question: Nickel(II) hydroxide (s) decomposes to produce nickel(II) oxide(s) and water.

Answer: Ni(OH)2(s) NiO(s) + H2O(l)

Write the balanced chemical equations for the following double-replacement reaction.

Question: Aqueous lithium iodide and aqueous silver nitrate react to produce solid silver iodide and aqueous lithium nitrate.

Answer: LiI(aq) + AgNO3(aq) AgI(s) + LiNO3(aq)

10.3 Reactions in Aqueous Solutions

In aqueous solutions, the solvent is always water. There are many possible solutes.

Many molecular compounds form ions when then they dissolve in water. When most ionic compounds dissolve in water, their ions separate.

When two aqueous solutions that contain ions as solutes are combined, the ions may react with one another. The solvent molecules do not usually react.

Reactions that occur in aqueous solutions are double-replacement reactions.

Three type of products produced during reactions in aqueous solutions are precipitates, water, and gases.

An ionic equation shows the details or reactions in aqueous solutions. A complete ionic equation shows all the particles in a solution as they exist. A net ionic equation includes only the particles that participate in a reaction in a solution.

Review Questions:

Write chemical, complete ionic, and net ionic equations for the following reactions.

Question: Aqueous solutions of potassium iodide and silver nitrate are mixed, forming the precipitates silver iodide.

Answer:Chemical equation:KI(aq) + AgNO3(aq) KNO3(aq) + AgI(s)

Complete ionic equation:K+(aq) + I-(aq) + Ag+(aq) + NO3

-(aq) K+(aq) + NO3-(aq) + AgI(s)

Net ionic equation:I-(aq) + Ag+(aq) AgI(s)

Question: Aqueous solutions of lithium sulfate and calcium nitrate are mixed, forming the precipitate calcium sulfate.

Answer:Chemical equation:Li2SO4(aq) + Ca(NO3)2(aq) 2LiNO3(aq) + CaSO4(s)

Complete ionic equation:2Li+(aq) + SO4

2-(aq) + Ca2+(aq) + 2NO3-(aq) 2Li+(aq) + 2NO3

-(aq) + CaSO4(s)

Net ionic equation:SO4

2-(aq) + Ca2+(aq) CaSO4(s)

Chapter 11 The Mole (ALBERT ALVAREZ, CARLOS QUINTANILLA, ANGELES PHAN, RUIKI BASILIO)

11.1 Measuring Matter

The Mole is a unit to count particles indirectly. One moles the amount of a pure substance that contains 6.02 x 10^23 representative particles.

One mole of carbon-12 atoms contains 12 grams of the isotope carbon-12.

Review Questions:

1. If you can compare 4.50 x 10^23 atoms Zn with 6.02 x 10^23, the number of the atoms in one mole, you can predict that the answer should be less than 10 moles.

Known: # of atoms = 4.50 x 10^23 atoms Zn, 1mole Zn = 6.02 x 10^23 atoms Zn Unknown: moles Zn = ? mole Answer: 7.48 mol Zn

2. Why is the mole an important unit to chemists?

11.2 Mass and the Mole

The molar mass of an element is the numerical equivalent of the atomic mass (a.m.u.) in grams.

The molar mass of any substance is the mass in grams of Avogadro’s number of representative particles of the substance.

Molar mass is used to convert from moles of an element to mass, and the inverse of molar mass is used to convert from mass of an element to moles.

Review Questions:

What is the mass in grams of 6.02 x 10^24 atoms Bi?Step 1: Divide by Avogadro’s numberAnswer: 2.09 x 10^3 g Bi

What is molar mass?

11.3 Moles of Compounds

Subscripts in a chemical formula indicate how many moles of each element are in one mole of the compound.

The Molar Mass of a compound is the sum of the masses of all the moles of elements present in the compound.

Review Questions:

What is the mass of 3.25 moles of sulfuric acid (H2 SO4)?Step 1: Find molar mass H2 SO4.Step 2: Make mole mass conversion.

Answer: 319 g H 2 SO 4

List three conversion factors used in molar conversions.

11.4 Empirical and Molecular Formulas (Ruiki Basilio & Angeles Phan)

The percent composition of a known compound can be calculated by dividing the mass of each element in one mole by the mass of a mole of the compound and multiplying by 100.

The subscripts in an empirical formulas are in a ratio of the smallest whole numbers of moles of the elements in the compound.

The molecular formula for a compound can be determined by finding the integer by which the mass of the empirical formula differs from the molar mass of the compound.

Review Questions:

1. Analysis of a chemical used in photographic developing fluid indicates a chemical composition of 65.45% C, 5.45% H, and 29.09% O. The molar mass is found to be 110.0g/mol. Determine the molecular formula.

2. An oxide of aluminum contains 0.545 g Al and 0.485 g O. Find the empirical formula for the oxide.

11.5 The Formula for a Hydrate

The formula for a hydrate consists of the formula for the ionic compound and the number of water molecules associated with one formula unit.

The name of a hydrate consists of the compound name followed by the word hydrate with a prefix indicating the number of water molecules associated with one mole of compound.

Anhydrous compounds are formed when hydrates are heated and the water of hydrated is driven off.

Review Questions:

1. A hydrate is found to have the following percent composition: 48.8% MgSO4 and 51.2% H2O. What is the formula for this hydrate?2. Explain how hydrates are named.

Chapter 15 Solutions (Lilian Orellana & Jocelyn Flores)

15.1 What are solutions?

A solute dissolves in a solvent during a process called solvation. When the solvent is water, the process also is called hydration.

Every substance has a characteristic solubility in a given solvent.

Factors that affect solubility include the nature of the solute and solvent, temperature, and pressure

Henry’s law states that the solubility (S) of a gas in a liquid is directly proportional to the pressure (P) of the gas above the liquid at a given temperature.

Practice Problems

1. If 0.55 g of a gas dissolves in 1.0 L of water at 20.0 kPa of pressure, how much will dissolve at 110.0 kPa of pressure?

S1 = 0.55 g = 0.55 g/L1.0 L

S2 = P2 x S 1 = 110.0 kPa x 0.55 g/L = 3.0 g/L P1 20.0 kPa

2. A gas has a solubility of 0.66 g/L at 10.0 atm of pressure. What is the pressure on a 1.0-L sample that contains 1.5 g of gas?

S2 = 1.5 g = 1.5 g/L1.0 L

P2 = S 2 x P1 = 1.5 g/L x 10.0 atm = 23 atm S1 0.66 g/L

15.2 Solution Concentration

The concentration of a solution is a quantitative measure of the amount of solute in a given amount of solvent or solution.

Measures of concentration include mass and volume, percentages, molarity, molality, and mole fraction.

A dilute solution can be prepared from a more concentrated standard stock solution.

Practice Problems:

1. What is the percent by mass of NaHCO3 in a solution containing 20 g NaHCO3 dissolved in 600 mL H2O?

600 mL H2O x 1.0 g/mL = 600 g H2O 20 g NaHCO3 X 100 = 3% 600 g H2O + 20 g Na HCO3

2. You have 1500.0 g of a bleach solution. The percent by mass of the solute sodium hypochlorite, NaOCl, is 3.62%. How many grams of NaOCl are in the solution?

3.62% = 100 x mass NaOCl1500.0 g

Mass NaOCl = 54.3 g

15.3 Colligative Properties of Solutions (Chanell Waddles & Monica Ajanel)

Physical properties affected by the concentration of the solute but not the nature of the solute are called colligative properties.

Colligative properties of solutions include vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.

15.4 Heterogeneous Mixtures

One of the key differences between solutions, colloids, and suspensions is particle size.

The random motion of colloidal dispersions due to molecular collisions is called Brownian motion.

The scattering of light by colloidal particles is called the Tyndall effect. The Tyndall effect can be used to distinguish colloids from solutions.

Equations for Chapter 15

Percent by mass = mass of solute/mass of solution X 100

Percent by Volume = volume of solute/volume of solution X 100

Molarity (M) = moles of solute/liters of solution

M 1V1 =M2 V2

Molality (m) = moles of solute/kilogram of solvent = moles of solute/1,000g of solvent

ΔTb =Kbm

Percent by mass = mass of solute x 100 mass of solution

Percent of Volume = volume of solute x100 volume of solution

Molarity (M) = moles of solute

liters of solution

Molality (m) = moles of solute = moles of solute kilogram of solvent 1000 g of solvent