Ch 14 Gas Law Problem

Assume that one cubic foot of air near a thermonuclear explosion is heated from zero degrees celsius to 546, 000 degrees celsius. what will the volume be?

Ch 22 Functional Groups

(40)Give the IUPAC name for each compound. Hint




22.1 Hydrocarbons

(37)Draw condensed structural formulas for pentane and hexane. Assume that the C ― H and C ― C bonds are understood. Hint

(38)Name the alkanes that have the following molecular or structural formulas.Hint

1.




CH3CH2CH3
2.


CH3(CH2)6CH3
3.


(39)Draw structures for the alkyl groups derived from methane, ethane, and propane. Hint


(41)Why are alkane molecules nonpolar? Hint
22.2 Unsaturated Hydrocarbons

(42)Give a systematic name for these alkenes. Hint

1.

CH3CH = CH2
2.

3.

4.

(43)Name and draw a structural formula for each alkene with the molecular formula C5H10. Hint
22.3 Isomers

(44)Draw and name all the structural isomers with the molecular formula C6H14. (You may wish to draw only the carbon skeletons.) Hint

(45)Draw one structural isomer of each compound. Hint

1.

2.

(46)Draw a structural formula or carbon skeleton for each of the following alkenes. If cis and trans forms are present, include both forms. Hint

1.

2-pentene
2.

2-methyl-2-pentene
3.

3-ethyl-2-pentene

(47)Do all molecules have optical isomers? Explain. Hint

(48)Can you draw a structural isomer of hexane (C6H14) that has an asymmetric carbon? Explain.Hint
22.4 Hydrocarbon Rings

(49)Draw a structural formula for each compound. Hint

1.

1,4-diethylbenzene
2.

2-methyl-3-phenylpentane
3.

1,3-dimethylbenzene

(50)Explain why both of these structures represent 1,2-diethylbenzene. Hint

22.5 Hydrocarbons From Earth’s Crust

(51)How are catalysts used in petroleum refining? Hint

(52)Rank these materials in order of increasing hardness: bituminous coal, peat, lignite, and anthracite coal.Hint

(53)What happens to the sulfur when coal burns?Hint
Understanding Concepts

(54)Why are the following names incorrect? What are the correct names? Hint

1.

2-dimethylpentane
2.

1,3-dimethylpropane
3.

3-methylbutane
4.

3,4-dimethylbutane

(55)For each hydrocarbon shown, identify the type of covalent bonds and name the compound.Hint

(56)Write structural formulas for these compounds.

1.

propyneHint
2.

cyclohexaneHint
3.

2-phenylpropaneHint
4.

2,2,4-trimethylpentaneHint

(57)Name the next three higher homologs of ethane. Hint

(58)Compare geometric and optical isomers.Hint

(59)Draw electron dot structures for each compound.

1.

etheneHint
2.

propane Hint
3.

ethyne Hint
4.

cyclobutane Hint

(60)Write an equation for the combustion of octane. Hint

(61)Compare these three molecular structures. Which would you expect to be most stable? Why? Hint

(62)The seven organic chemicals produced in the largest amounts in the United States in a recent year are listed in the table below. Answer the following questions based on the data given. Hint

1.

How many billion kilograms of aromatic compounds were produced?
2.

Of the total mass of all seven compounds produced, what percent by mass was made up of aliphatic compounds?

(63)Are these two structures geometric isomers? Explain your answer. Hint

(64)Use the labeled features in the molecular structure to answer the following questions.

Ch 13 Kinetic Theory

Use the graph to answer Questions 1 and 2.



(1)What is the normal boiling point of ethanol?Hint

(2)Can chloroform be heated to 90°C in an open container?Hint

(3)Which sequence has the states of CH3OH correctly ordered in terms of increasing average kinetic energy?Hint

1. CH3OH(s), CH3OH(g), CH3OH(l)
2. CH3OH(g), CH3OH(l), CH3OH(s)
3. CH3OH(l), CH3OH(g), CH3OH(s)
4. CH3OH(s), CH3OH(l), CH3OH(g)

Use the drawing to answer Questions 4–6. The same liquid is in each flask.

(4)In which flask is the vapor pressure lower? Give a reason for your answer.Hint

(5)In which flask is the liquid at the higher temperature? Explain your answer.Hint

(6)How can the vapor pressure in each flask be determined?Hint

For each question there are two statements. Decide whether each statement is true or false. Then decide whether Statement II is a correct explanation for Statement I.

(7)In an open container, the rate of evaporation of a liquid always equals the rate of condensation.

BECAUSE

A dynamic equilibrium exists between the liquid and its vapor in an open container. Hint

(8)Water boils at a temperature below 100°C on top of a mountain.

BECAUSE

Atmospheric pressure decreases with an increase in altitude. Hint

(9)The temperature of a substance always increases as heat is added to the substance.

BECAUSE

The average kinetic energy of the particles in a substance increase with an increase in temperature.Hint

(10)Solids have a fixed volume.

BECAUSE

Particles in a solid cannot move.Hint

(11)Gases are more compressible than liquids.

BECAUSE

There is more space between particles in a gas than between particles in a liquid.Hint

Ch 14 Gas Laws 2

(1)A gas in a balloon at constant pressure has a volume of 120.0 mL at −123°C. What is its volume at 27.0°C?Hint

1. 60.0 mL
2. 240.0 mL
3. 26.5 mL
4. 546 mL

(2)If the Kelvin temperature of a gas is tripled and the volume is doubled, the new pressure will beHint

1. 1/6 the original pressure.
2. 2/3 the original pressure.
3. 3/2 the original pressure.
4. 5 times the original pressure.

(3)Which of these gases effuses fastest?Hint

1. Cl2
2. NO2
3. NH3
4. N2

(4)All the oxygen gas from a 10.0-L container at a pressure of 202 kPa is added to a 20.0-L container of hydrogen at a pressure of 505 kPa. After the transfer, what are the partial pressures of oxygen and hydrogen?Hint

1. Oxygen is 101 kPa; hydrogen is 505 kPa.
2. Oxygen is 202 kPa; hydrogen is 505 kPa.
3. Oxygen is 101 kPa; hydrogen is 253 kPa.
4. Oxygen is 202 kPa; hydrogen is 253 kPa.

(5)Which of the following changes would increase the pressure of a gas in a closed container?Hint

1.

Part of the gas is removed.
2.

The container size is decreased.
3.

Temperature is increased.

1. I and II only
2. II and III only
3. I and III only
4. I, II, and III

(6)A real gas behaves most nearly like an ideal gasHint

1. at high pressure and low temperature.
2. at low pressure and high temperature.
3. at low pressure and low temperature.
4. at high pressure and high temperature.

Use the graphs to answer Questions 7–10. A graph may be used once, more than once, or not at all.

Which graph shows each of the following?

(7)directly proportional relationshipHint Link

(8)graph with slope = 0Hint Link

(9)inversely proportional relationshipHint Link

(10)graph with a constant slopeHint Link

Use the drawing to answer Questions 11 and 12.

(11)Bulb A and bulb C contain different gases. Bulb B contains no gas. If the valves between the bulbs are opened, how will the particles of gas be distributed when the system reaches equilibrium? Assume none of the particles are in the tubes that connect the bulbs.Hint

(12)Make a three-bulb drawing with 6 blue spheres in bulb A, 9 green spheres in bulb B, and 12 red spheres in bulb C. Then draw the setup to represent the distribution of gases after the valves are opened and the system reaches equilibrium.Hint

Ch 14 Gas Laws

How does kinetic theory explain the compressibility of gases?Hint

Use this description to answer Questions 65 and 66.

A teacher adds 1 mL of water to an empty metal soda can. The teacher heats the can over a burner until the water boils and then quickly plunges the can upside down in an ice-water bath. The can immediately collapses inward as though crushed in a trash compactor.

(65)Use kinetic theory to explain why the can collapsed inward.Hint

(66)If the experiment were done with a dry can, would the results be similar? Explain.Hint

Bonus Question Ch 19 (EP 6)

Find Kb and pKb for the acetate ion CH3COO-. The ionization constant of CH3COOH is Ka=1.75 x 10-5 (this is to -5 power); Kw=1.00 x 10-14--this is to the -14 power.

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Bonus Question Set 1 Ch 2

Explain why mass cannot be used as a property to identify a sample of matter?

On page 59 of your Prentice Hall Chemistry Book, answer questions 60-63

(28) Key Concept How does a chemical change affect the composition of matter? Hint

(29) Key Concept Name four possible clues that a chemical change has taken place.Hint

(30) Key Concept In a chemical reaction, how does the mass of the reactants compare with the mass of the products?Hint

(31)What is the main difference between physical changes and chemical changes?

Bonus Question Set 1 ch 1

You are asked to design an experiment to answer the question which is the best paper towel in thailand?

a) what is the manipulated variable in your experiment
b) list three possible responding variables that could be used to define best
c) pick one of the responding variables and rewrite the question as a hypothesis
d) list at least five factors that must be kept constant when you test the hypothesis

Patterns of Chemical Reactivity

Patterns of Chemical Reactivity

Using the Periodic Table

• As a consequence of the good ordering of the periodic table, the properties of compounds of elements vary in a systematic manner.
• Example: All the alkali metals (M) react with water as follows:

2 M (s) + 2 H2O (l) --> 2 MOH (aq) + H2 (g)

 The reactions become more vigorous as we move from Li to Cs
 Sodium reacts with water to produce an orange flame.
 Potassium reacts with water to produce a blue flame.
 The reaction of potassium with water produces so much heat that the hydrogen gas produced usually ignites with a loud pop.


Combustion in Air

• Combustion reactions are rapid reactions that produce a flame.
 Combustion is the burning of a substance in air.
 Example: Propane combusts to produce carbon dioxide and water:

C3H8 (g) + 5 O2 (g) --> 3 CO2 (g) + 4 H2O (l)


Combination and Decomposition Reactions

• In combination reactions two or more substances react to form one product.
• Combination reactions have more reactants than products.
 Consider the reaction:

2 Mg (s) + O2 (g) --> 2 MgO (s)

• Since there are fewer products than reactants, the Mg has combined with O2 to form MgO.
• Note that the structure of the reactants has changed:
 Mg consists of closely packed atoms, and O2 consists of dispersed molecules.
 MgO consists of a lattice of Mg2+ and O2- ions.

• In decomposition reactions one substance undergoes a reaction to produce two or more other substances.
• Decomposition reactions have more products than reactants.
 Consider the reaction that occurs in an automobile air bag:
2 NaN3 (s) --> 2 Na (s) + 3 N2 (g)

 Since there are more products than reactants, the sodium azide has decomposed into Na metal and N2 gas.

Chemical Equations

Chemical Equations

• Lavoisier observed that mass is conserved in a chemical reaction.
 This observation is known as the law of conservation of mass.
• The quantitative nature of chemical formulas and reactions is called stoichiometry.
• Chemical equations give a description of a chemical reaction.
• There are two parts to any equation:
 Reactants (written to the left of the arrow) and
 Products (written to the right of the arrow):

2 H2 + O2 --> 2 H2O
• There are two sets of numbers in a chemical equation:
 Numbers in front of the chemical formulas (called stoichiometric coefficients) and
 Numbers in the formulas (they appear as subscripts).
• Stoichiometric coefficients give the ratio in which the reactants and products exist.
• The subscripts give the ratio in which the atoms are found in the molecule.
 Example:
• H2O means there are two H atoms for each one molecule of water.
• 2 H2O means that there are two water molecules present.
• Note: In 2 H2O there are four hydrogen atoms present (two for each water molecule).
• Matter cannot be lost in chemical reactions.
 Therefore, the products of a chemical reaction have to account for all the atoms present in the reactants.
• Consider the reaction of methane with oxygen.

CH4 + O2 --> CO2 + H2O

• Counting atoms in the reactants:
 1 C;
 4 H; and
 2 O
• In the products:
 1 C;
 2H; and
 3O
• It appears as though H has been lost and C has been created.
• To balance the equation, we adjust the stoichiometric coefficients:

CH4 + 2 O2 --> CO2 + 2 H2O

Periodic Table More

Excited Atoms and the Fourth of July
1. What is light, and how do various colors of light differ?
2. What is going on at the level of atoms and molecules when fireworks produce colored light?
3. How does the instability of copper chloride at high temperatures ineterfere with its ability to emit blue
light?


3-1 How are the Elements Organized?

Objectives:
* Describe the organization of the modern periodic table.
* Use the periodic table to obtain information abour the properties of elements..
* Explain how the names and symbols of elements are derived.
* Identify common metals, nonmetals, and metalloids, and noble gases.

PROPERTIES OF ELEMENTS RELATE TO ATOMIC STRUCTURES

The periodic table shows all the elements




The periodic table is organized by properties



Group - a series of elements that form a column in the periodic table.
Period - a series of elements that form a horizontal row in the periodic table.

REGIONS OF THE PERIODIC TABLE

The periodic table contains regions of similar elements

Metals -
Nonmetals -
Metalloids -
Noble Gases -

Metals for the largest region of the table
Examples:

Nonmetals are the second largest region of the table
Examples:


Metalloids have properties of both metals and nonmetals
Examples:


Noble gases are on the far right of the periodic table
Examples:


Essential Elements - elements needed for health
Examples:

Naming Inorganic Compounds

Naming Inorganic Compounds

Names and Formulas of Ionic Compounds
• Chemical nomenclature – naming of substances.
• Divided into organic compounds (those containing C, usually in combination with H, O, N, or S) and inorganic compounds (all other compounds)

1. Positive Ions (Cations)
• Cations formed from a metal have the same name as the metal
 Example: Na1+ = sodium ion
• If the metal can form more than one cation, then the charge is indicated in parentheses in the name.
 Examples: Cu1+ = copper (I); Cu2+ = copper (II) [Stock system]
• Cations formed from nonmetals end in –ium.
 Example: NH41+ ammonium ion

2. Negative Ions (anions)
• Monatomic anions (with only one atom) use the ending –ide.
Example: Cl1- is chloride ion

• Some simple polyatomic anions also have the ending -ide; hydroxide, cyanide, and peroxide ions.
• Polyatomic ions (with many atoms) containing oxygen are called oxyanions.
 These end in –ate or –ite. (The one with more oxygen uses –ate.)
 Examples: NO31- is nitrate, NO21- is nitrite.

• Polyatomic anions containing oxygen with additional hydrogens are named by adding hydrogen or bi (one H), dihydrogen (two H), etc. to the name, as follows:
 CO32- is the carbonate anion
 HCO3 is the hydrogen carbonate (or bicarbonate) anion
 H2PO4 is the dihydrogen phosphate anion


3. Ionic Compounds
• These are named cation and anion
 Example: BaBr2 = barium bromide


Names and Formulas of Acids

• The names of acids are related to the names of anions:
 -ide becomes hydro- ...-ic acid; Examples: HCl hydrochloric acid
 -ate becomes –ic acid; HClO4 perchloric acid
 -ite becomes –ous acid. HClO hypochlorous acid


Names and Formulas for binary Molecular Compounds

• Binary molecular compounds contain two elements.
• The most metallic element (i.e., the one farthest to the left on the periodic table) is usually written first. Exception: NH3 (ammonia)
• If both elements are in the same group, the lower one is written first.
• Greek prefixes are used to indicate the number of atoms (e.g. mono, di, tri).
 The prefix mono is never used with the first element (i.e., carbon monoxide, CO).
 Examples:
• Cl2O dichlorine monoxide
• N2O4 dinitrogen tetroxide
• NF3 nitrogen trifluoride
• P4S10 tetraphosphorus decasulfide



CHEMISTRY The Central Science 8th Edition Brown, LeMay, Bursten Ch 2: Atoms, Molecules, and Ions

Compounds Ionic Covalent

2.5 Molecules and Molecular Compounds

• A molecule consists of two or more atoms bound together.
• Each molecule has a chemical formula.
• The chemical formula indicates
1. Which atoms are found in the molecule, and
2. In what proportion they are found
• Compounds composed of molecules are molecular compounds
• These contain at least two types of atoms.
• Different forms of an element have different chemical formulas are known as allotropes. Allotropes differ in their chemical and physical properties. (Chapter 7 will give more information on allotropes of common elements)

Molecular and Empirical Formulas
• Molecular formulas
 Give the actual numbers and types of atoms in a molecule.
 Examples: H2O, CO2, CO, CH4, H2O2, O2, O3, and C2H4
• Empirical formulas
 Give the relative numbers and types of atoms in a molecule (they give the lowest whole-number ratio of atoms in a molecule).
 Examples: H2O, CO2, CO, CH4, HO, CH2

Picturing Molecules
• Molecules occupy three-dimensional space.
• However, we often represent them in two dimensions.
• The structural formula gives the connectivity between individual atoms in the molecule.
• The structural formula may or may not show the three-dimensional shape of the molecule.
• If the structural formula does not show the shape of the molecule, then either a perspective drawing, ball-and-stick model, or space-filling model is used.
 Perspective drawings use dashed lines and wedges to represent bonds receding and emerging from the plane of the paper.
 Ball-and-stick models show atoms as contracted spheres and the bonds as sticks. The angles in the ball-and-stick model are accurate.
 Space-filling models give an accurate representation of the relative sizes of the atoms and the 3D shape of the molecule.

2.6 Ions and Ionic Compounds

• If electrons are added or removed from a neutral atom, an ion is formed.
• When an atom or molecule loses electrons, it becomes positively charged.
 Positively charged ions are called cations.
• When an atom or molecule gains electrons, it becomes negatively charged.
 Negatively charged ions are called anions.
• In general, metal atoms tend to lose electrons, and nonmetal atoms gain electrons.
• When molecules lose electrons, polyatomic ions are formed (e.g. SO42-, NO31-)


Predicting Ionic Charges

• An atom or molecule can lose more than one electron.
• Many atoms gain or lose enough electrons to have the same number of electrons as the nearest noble gas (group 18 or 8A).
• The number of electrons an atom loses is related to its position on the periodic table.

Ionic Compounds

• A great deal of chemistry involves the transfer of electrons between species.
• Example:
 To form NaCl. the neutral sodium atom, Na, must lose an electron to become a cation: Na1+.
 The electron cannot be lost entirely, so it is transferred to a chlorine atom, Cl, which then becomes an anion: Cl1-.
 The Na1+ and Cl1- ions are attracted to form an ionic NaCl lattice, which crystallizes.
• NaCl is an example of an ionic compound – consisting of positively charged cations and negatively charged anions.
 Important: note that there are no easily identified NaCl molecules in the ionic lattice. Therefore we cannot use molecular formulas to describe ionic substances.
• In general, ionic compounds are usually combinations of metals and nonmetals, whereas molecular compounds are generally composed of nonmetals only.
• Writing empirical formulas for ionic compounds:
 You need to know the ions of which it is composed.
 The formula must reflect the electrical neutrality of the compound.
 You must combine cations and anions in a ratio so that the total positive charge is equal to the total negative charge.
 Example: consider the formation of Mg3N2:
• Mg loses two electrons to become Mg2+
• Nitrogen gains three electrons to become N3-
• For a neutral species, the number of electrons lost and gained must be equal.
• However, Mg can lose electrons only in twos, and N can accept electrons only in threes.
• Therefore, Mg needs to form 3 Mg2+ ions (total 3 x 2 positive charges) and 2 N atoms need to form 2 N3- ions (total 2 x 3 negative charges)
• Therefore, the formula is Mg3N2
•

Chemistry and Life: Elements Required by Living Organisms

• Of the 112 elements known, only about 26 are required for life.
• Water accounts for more than 70 percent of the mass of the cell.
• Carbon is the most common solid constituent of cells.
• The most important elements for life are H, C, N, O, P, and S (red).
• The next most important ions are Na1+, Mg2+, K1+, Ca2+, and Cl1- (blue)
• The other 15 elements are needed only in trace amounts (green).



CHEMISTRY The Central Science 8th Edition Brown, LeMay, Bursten Ch 2: Atoms, Molecules, and Ions

Periodic Table

2.4 The Periodic Table

• The Periodic Table is used to organize the elements in a meaningful way.
• As a consequence of this organization, there are periodic properties associated with the periodic table.
• Columns in the periodic table are called groups.
 Several numbering conventions are used (i.e. groups may be numbered from 1 to 18, or from 1A to 8A, and from 1B to 8B).
• Rows in the periodic table are called periods.
• Some of the groups in the periodic table are given special names.
 These names indicate the similarities among group members.
 Examples:
• Group 1 or (1A): alkali metals
• Group 17 or (7A): halogens
• Metallic elements are located on the left hand-side of the periodic table (most of the elements are metals).
• Nonmetallic elements are located in the top right-hand side of the periodic table.
• Elements with properties similar to both metals and nonmetals are called metalloids and are located at the interface between the metals and nonmetals.
 These include the elements B, Si, Ge, As, Sb, and Te.
• Metals tend to be malleable, ductile, and lustrous and are good thermal and electrical conductors. Nonmetals generally lack these properties; they tend to be brittle solids, dull in appearance, and do not conduct heat of electricity well.

CHEMISTRY The Central Science 8th Edition Brown, LeMay, Bursten Ch 2: Atoms, Molecules, and Ions

Atomic Theories of Matter

2.1 The Atomic Theory of Matter

• Greek philosophers: Can matter be subdivided into fundamental particles?
• Democritus (460 – 370 BC): All matter can be divided into indivisible atomos.
• Dalton: Proposed atomic theory with the following postulates:
 Elements are composed of atoms.
 All atoms of an element are identical.
 In chemical reactions atoms are not changed into different types of atoms. Atoms are neither created not destroyed.
 Compounds are formed when atoms of elements combine.
• Atoms are the building blocks of matter.
• Law of constant composition: the relative kinds and numbers of atoms are constant for a given compound.
• Law of conservation of mass: during a chemical reaction, the total mass before reaction is equal to the total mass after reaction.
 Conservation means something can neither be created nor destroyed. Here it applies to matter (mass). Later we will apply it to energy (Chapter 5).

• Law of multiple proportions: if two elements A and B combine to form more than one compound, then the mass of B that combines with the mass of A is a ratio of small whole numbers.
• Dalton’s theory predicted the law of multiple proportions.


2.2 The Discovery of Atomic Structure

• By 1850 scientists knew that atoms were composed of charged particles.
• Subatomic particles: those particles that make up the atom,
• Recall the law of electrostatic attraction: like charges repel and opposite charges attract.

Cathode Rays and Electrons
• Cathode rays were first discovered in the mid- 1800s from studies of electrical discharge through partially evacuated tubes (cathode ray tubes or CRTs).
 Computer terminals were once popularly referred to as CRTs (cathode ray tubes).
 They are now commonly called VDTs (video display terminals)
• Cathode rays = radiation produced when high voltage is applied across the tube.
• The voltage causes negative particles to move from the negative electrode (cathode) to the positive electrode (anode).
• The path of the electrons can be altered by the presence of a magnetic field.
• Consider cathode rays leaving the positive electrode through a small hole.
 If they interact with a magnetic field perpendicular to an applied electric field, then the cathode rays can be deflected by different amounts.
 The amount of deflection also depends on the applied magnetic and electric fields.
 The amount of deflection also depends on the charge-to-mass ratio of the electron.
 In 1897 Thomson determined the charge-to-mass ratio of an electron.
• Charge-to-mass ratio: 1.76 x 108 C / g
• C is a symbol for coulomb
o SI unit of electric charge
• Millikan Oil-Drop Experiment
 Goal: find the charge on the electron to determine its mass.
 Oil drops were sprayed above a positively charged plate containing a small hole.
 As the oil drops fall through the hole they acquire a negative charge.
 Gravity forces the drops downward. The applied electric field forces the drops upward.
 When a drop is perfectly balanced, then the weight of the drop is equal to the electrostatic force of attraction between the drop and the positive plate.
 Millikan carried out the above experiment and determined the charges on the oil drops to be multiples of 1.60 x 10-19 C.
 He concluded that the charge on the electron must be 1.60 x 10-19 C.
• Knowing the charge-to-mass ratio of the electron, we can calculate the mass of the electron:

Radioactivity
• Radioactivity is the spontaneous emission of radiation.
• Consider the following experiment:
 A radioactive substance is placed in a lead shield containing a small hole so that a beam of radiation is emitted from the shield.
 The radiation is passed between two electrically charged plates and detected.
 Three spots are observed on the detector:
• A spot deflected in the direction of the positive plate.
• A spot that is not affected by the electric field.
• A spot deflected in the direction of the negative plate.
 A large deflection toward the positive plate corresponds to radiation that is negatively charged and of low mass. This is called -radiation (consists of electrons).
 No deflection corresponds to neutral radiation. This is called -radiation (similar to X-rays).
 A small deflection toward the negatively charged plate corresponds to high mass, positively charged radiation. This is called -radiation (positively charged core of a helium atom).
 X-rays and  radiation are true electromagnetic radiation, whereas - and -radiation are actually streams of particles – helium nuclei and electrons, respectively.

The Nuclear Atom
• The ‘plum pudding’ model: an early picture of the atom.
• The Thomson model pictures the atom as a sphere with small electrons embedded in a positively charged mass.
• Rutherford carried out the following experiment:
 A source of -particles was placed at the mouth of a circular detector.
• The -particles were shot through a piece of gold foil.
• Both the gold nucleus and the a-particle are positively charged, so they repel each other.
• Most of the -particles went straight through the foil without deflection.
• If the Thomson model of the atom was correct, then Rutherford’s results was impossible.
• Rutherford modified Thomson’s model as follows:
 Assume that the atom is spherical, but the positive charge must be located at the center.
 In order for the majority of -particles shot through a piece of foil to be undeflected, the majority of the atom must consist of empty space where the electrons can be found.
 To account for the small number of large deflections of the -particles, the center or nucleus of the atom must consist of a dense positive charge.



2.3 The Modern View of Atomic Structure

• The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons).
• Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus.
• Electrons are located outside of the nucleus. Most of the volume of the atom is the space where electrons are found.
• The quantity 1.602 x 10-19 C is called electronic charge. The charge on an electron is –1.602 x 10-19 C; the charge on a proton is +1.602 x 10-19 C.
• Masses are so small that we define the atomic mass unit, amu.
 1 amu = 1.66054 x 10-24 g
 The mass of a proton is 1.0073 amu, a neutron is 1.0087 amu, and an electron is 5.486 x 10-4 amu.
 The angstrom is a convenient non-SI unit of length used to denote atomic dimensions.
• Since most atoms have radii around 1 x 10-10 m, we define 1  = 1 x 10-10 m.

Isotopes, Atomic Numbers, and Mass Numbers
• Atomic number (Z) = number of protons in the nucleus
• Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons).
• By convention, for element X, we write
• Isotopes have the same Z but different A.
 There can be a variable number of neutrons for the same number of protons. Isotopes have the same number of protons but different numbers of neutrons.
 An atom of a specific isotope is called a nuclide.
 Examples: Nuclides of hydrogen include:
• H-1 (protium); H-2 (deuterium); H-3 (tritium): tritium is radioactive

CHEMISTRY The Central Science 8th Edition Brown, LeMay, Bursten Ch 2: Atoms, Molecules, and Ions

Properties of Matter and Derived Units

1.3 Properties of Matter

Physical vs. Chemical Properties

· Physical properties: Measured without changing the substance (e.g. color, density, odor, melting point).

· Chemical properties: Describe how substances react or change to form different substances (e.g. hydrogen burns in oxygen).

· Intensive properties: Do not depend on the amount of substance present (e.g. temperature, melting point); give an idea of the composition of a substance.

· Extensive properties: Depend on quantity of substance present (e.g. mass, volume).

Physical and Chemical Change

· Physical change: Substance changes physical appearance without altering its identity (e.g. changes of state).

· Chemical changes (or chemical reactions): Substances transform into chemically different substances (i.e. identity changes, e.g. decomposition of water, explosion of nitrogen triiodide).

The Scientific Method

The scientific method: guidelines for the practice of science.

· Collect data (observe, experiment, etc.)

· Look for patterns, try to explain them and develop a hypothesis.

· Test hypothesis; refine it.

· Bring all information together into a scientific law (concise statement or equation that summarizes tested hypotheses).

· Bring hypotheses and laws together into a theory. a theory should explain general principles.

1.4 Units of Measurement

· Many properties of matter are quantitative.

· A measured quantity must have BOTH a number and a unit.

· The units most often used for scientific measurement are those of the metric system.

SI Units

· 1960: all scientific units use Systeme International d’Unites (SI Units)

· There are seven base units.

· Smaller and larger units are decimal fractions or multiples of the base units.

Length and Mass

· SI base unit of length = meter (1 m = 1.0936 yards).

· SI base unit of mass (not weight) = kilogram (1 kg = 2.2 pounds).

§ Mass is a measure of the amount of material in an object.

Temperature

· Scientific studies use Celsius and Kelvin scales.

· Celsius scale: Water freezes at 0 oC and boils at 100 oC (sea level).

· Kelvin scale (SI unit)

§ Water freezes at 273.15 K and boils at 373.15 K (sea level).

§ Based on properties of gases.

§ Zero is lowest possible temperature (absolute zero).

§ 0 K = -273.15 oC.

· Fahrenheit (not used in science)

§ Water freezes at 32 oF and boils at 212 oF (sea level)

§ Conversions:

oF - 32 = 1.8 oC K = oC + 273.15

Derived SI Units

· These are formed from the seven base units.

· Example: velocity is distance traveled per unit time, so units of velocity are units of distance (m) divided by units of time (s): m / s

Volume

· Units of volume = (units of length)3 e.g. m3

· This unit is unrealistically large, so we use more reasonable units:

§ cm3 (also known as mL or cc (cubic centimeters))

§ dm3 (also known as liters, L)

§ Important: the liter is not an SI unit.

Density

· Used to characterize substances.

· Density is defined as mass divided by volume.

· Units are usually g / cm 3.

· Originally based on mass (the density was defined as the mass of 1.00 g of pure water).

Significant Figures--Conversions

1.5 Uncertainty in Measurement

· Two types of numbers:

· Exact numbers (known by counting or definition)

· Inexact numbers (derived from measurement)

· All measurements have some degree of uncertainty or error associated with them.

Precision and Accuracy

· Precision: how well measured quantities agree with one another.

· Accuracy: how well measured quantities agree with the “true value.”

· Figure 1.25 is very helpful in making this distinction.

Significant Figures

· In a measurement it is useful to indicate the exactness of the measurement. This exactness is reflected in the number of significant figures.

· Guidelines for determining the number of significant figures in a measured quantity:

§ The number of significant figures is the number of digits known with certainty plus one uncertain digit. (Example 2.2405 g means we are sure that the mass is 2.240 g, but we are uncertain about the nearest 0.0001 g.)

· Final calculations are only as significant as the least significant measurement.

· Rules:

1. Nonzero numbers are always significant.

2. Zeros between nonzero numbers are always significant.

3. Zeros before the first nonzero digit are not significant. (Example: 0.0003 has one significant figure.)

4. Zeros at the end of the number after a decimal place are significant.

5. Zeros at the end of the number after a decimal place are ambiguous (e.g. 10,300 g).

· Method:

1) Write the number in scientific notation.

2) The number of digits remaining is the number of significant figures.

3) Examples:

· 2.50 x 104 cm has 3 significant figures as written

· 1.03 x 104 g has 3 significant figures

· 1.030 x 104 g has 4 significant figures

· 1.0300 x 104 has 5 significant figures

Significant Figures in Calculations

· Multiplication and Division

§ Report to the least number of significant figures

· (e.g. 6.221 cm x 5.2 cm = 32 cm2)

· Addition and Subtraction

§ Report to the least number of decimal places

· (e.g. 20.4 g - 1.322 g = 19.1 g).

· In multiple-step calculations always retain an extra significant figure until the end to prevent rounding errors.

1.6 Dimensional Analysis

· Method of calculation utilizing a knowledge of units.

· Given units can be multiplied and divided to give the desired units.

· Conversion factors are used to manipulate units:

§ Desired unit = given unit x (conversion factor)

· The conversion factors are simple ratios:

§ Conversion factor = (desired unit) / (given unit)

Using Two or More Conversion Factors

· We often need to use more than one conversion factor in order to complete a problem.

· When identical units are found in the numerator and denominator of a conversion. they will cancel. The final answer MUST have the correct units.

§ Example:

· Suppose that we want to convert length in meters to length in inches. We can do this conversion with the following conversion factors:

1 meter = 100 centimeters and 1 inch = 2.54 centimeters

· The calculation will involve both conversion factors; the units of the final answer will be inches:

Conversions Involving Volume

· We often will encounter conversions from one measure to a different measure.

§ Example:

1. Suppose that we wish to know the mass in grams of 2.00 cubic inches of gold given that the density of the gold is 19.3 g/cm .

2. We can do this conversion with the following conversion factors:

2.54 cm = 1 inch and 1 cm3 = 19.3 g gold

3. The calculation will involve both of these factors:

x g gold =

4. Note that the calculation will NOT be correct unless the centimeter to inch conversion is cubed! Both the units AND the number must be cubed.

Summary of Dimensional Analysis

· In dimensional analysis always ask three questions:

1. What data are we given?

2. What quantity do we need?

3. What conversion factors are available to take us from what we are given to what we need?

from CHEMISTRY The Central Science 8th Edition Brown, LeMay, Bursten Ch 1: Introduction: Matter and

Classification and Introduction

Chemistry: Introduction: Matter and Measurement Lecture Outline

1.1 The Study of Chemistry

· Chemistry – study of properties of materials and changes that they undergo.

§ Can be applied to all aspects of life (e.g. development of pharmaceuticals, leaf color change in fall).

The Molecular Perspective of Chemistry

Chemistry involves the study of the properties and behavior of matter.

· Matter

§ Physical material of the universe

§ Has mass

§ Occupies space

§ ~100 Elements constitute all matter

§ Elements:

· Made up of unique atoms.

· Names of the elements are derived from a wide variety of sources

(e.g. Latin or Greek, mythological characters, names of people or places).

§ Molecules:

· Combinations of atoms held together in specific shapes.

· Macroscopic (observable) properties of matter relate to microscopic realms of atoms.

· Properties relate to composition (types of atoms present) and structure (arrangement of atoms) present.

Why Study Chemistry?

We study chemistry because:

· It has a considerable impact on society (health care, food, clothing, conservation of natural resources, environmental issues, etc.).

· It is part of your curriculum! Chemistry serves biology, engineering, agriculture, geology, physics, etc. Chemistry is the central science.

1.2 Classifications of Matter

Matter is classified by state (solid, liquid, or gas) or by composition (element, compound, or mixture).

States of Matter

On the macroscopic level:

· Gas No fixed volume or shape, conforms to volume and shape of container, compressible.

· Liquid Volume independent of container, no fixed shape, incompressible.

· Solid Volume and shape independent of container, rigid, incompressible.

On the molecular level

· Gas Molecules far apart, move at high speeds, collide often.

· Liquid Molecules closer than those in gas, move rapidly but can slide over one another.

· Solid Molecules packed closely in definite arrangements.

Pure Substances and Mixtures

· Pure Substance

§ Matter with fixed composition and distinct proportions.

§ Elements (cannot be decomposed into simpler substances, i.e. only one kind of atom) or compounds (consist of two or more elements).

· Mixtures

§ Combination of two or more pure substances.

§ Variable composition.

§ Heterogeneous (do not have uniform composition, properties and appearance, e.g. sand).

§ Homogeneous (uniform throughout, e.g. air). Homogeneous mixtures are solutions.

Separation of Mixtures

Key: Separation techniques exploit differences in properties of the components.

· Filtration: Remove solid from liquid

· Distillation: Boil off one or more components of the mixture

· Chromatography: Exploit solubility of components

Elements

· 112 known

· Vary in abundance

· Each is given a unique name

· Organized in periodic table

· Each given a one- or two-letter symbol derived from its name.

Compounds

· Combination of elements

Example: The compound H₂O is a combination of the elements H and O.

· The opposite of compound formation is decomposition

· Compounds have different properties than their component elements (e.g. water is liquid, hydrogen and oxygen are both gases at the same temperature and pressure).

· Law of constant (definite) proportions (Proust): A compound always consists of the same combination of elements (e.g. water is always 11 percent H and 89 percent O).

Measurement