Ch 100: Fundamentals for Chemistry

Ch 100: Fundamentals for Chemistry
Chapter 1: Introduction Lecture Notes

What is Chemistry? Chemistry is often described as the “central” science Chemistry is the study of matter Matter is the “stuff” that makes up the universe, i.e. anything that has mass and occupies space The fundamental questions of Chemistry are: How can matter be described? How does one type of matter interact with other types of matter? How does matter transform into other forms of matter?

Major Developments in Chemistry I
~400 BC: Democritus proposed the concept of the “atom” ~300 BC: Aristotle developed 1st comprehensive model of matter ~700 AD: Chinese alchemists invent gunpowder 1661: Robert Boyle proposed the concept of elements : Lavoisier proposed the concept of compounds & the Law of Mass Conservation 1774: Priestly isolates oxygen 1797: Proust proposed the Law of Definite Proportions 1803: Dalton re-introduces the concept of the atom and establishes Dalton’s Laws 1869: Mendeleev creates the 1st Periodic Table 1910: Rutherford proposes the “nuclear” model of the atom 1915: Bohr proposes a “planetary” model of the hydrogen atom 1920: Schroedinger publishes his wave equation for hydrogen 1969: Murray Gell-Mann proposes the theory of QCD (proposing the existence of quarks)

Major Developments in Chemistry II
Discovery of subatomic particles: 1886: Proton (first observed by Eugene Goldstein) 1897: Electron (JJ Thompson) 1920: Proton (named by Ernest Rutherford) 1932: Neutron (James Chadwick) Other Important Discoveries: 1896: Antoine Henri Becquerel discovers radioactivity 1911: H. Kamerlingh Onnes discovers superconductivity in low temperature mercury 1947: William Shockley and colleagues invent the first transistor 1996: Cornell, Wieman, and Ketterle observe the 5th state of matter (the Bose-Einstein condensate) in the laboratory

Scientific Method 1. (OBSERVATION) Recognize a problem
Make observation Formulate a question 2. (EXPLANATION) Make an educated guess - a hypothesis Predict the consequences of the hypothesis 3. (VALIDATION) Perform experiments to test the predictions Does experimental data support or dispute hypothesis? 4. Formulate the simplest rule that organizes the 3 main ingredients - develop a theory

EXPLANATIONS

Bottom Line: The Scientific Attitude
All hypotheses must be testable (i.e. there must be a way to prove them wrong!!) Scientific: “Matter is made up of tiny particles called atoms” Non-Scientific: “There are tiny particles of matter in the universe that will never be detected”

The Particulate Nature of Matter
Matter is the tangible substance of nature, anything with mass that occupies space At the most fundamental level, matter is discrete or particulate in nature The smallest, most basic units of matter are called atoms All matter is thus comprised of individual atoms, or specific combinations of atoms called molecules Molecules can be broken apart into their constituent atoms but atoms cannot be further broken apart and still retain the properties of matter Matter can exist in one or more physical states (or phases)

Solid → Liquid → Gas Solid ← Liquid ← Gas States of Matter State Shape
+Energy +Energy State Shape Volume Compress Flow Solid Keeps Keeps No No Shape Volume Liquid Takes Keeps No Yes Shape of Volume Container Gas Takes Takes Yes Yes Shape of Volume of Container Container Solid ← Liquid ← Gas +Energy +Energy 7

Classification of Matter
Matter can be classified as either Pure or Impure: Pure Element: composed of only one type of atom Composed of either individual atoms or molecules (e.g. O2) Compound: composed of more than one type of atom Consists of molecules Impure (or mixture) Homogeneous: uniform throughout, appears to be one thing Pure substances Solutions (single phase homogeneous mixtures) Suspensions (multi-phase homogeneous mixtures) Heterogeneous: non-uniform, contains regions with different properties than other regions Pure Substance Constant Composition Homogeneous Mixture Variable Composition Matter

Separation of Matter A pure substance cannot be broken down into its component substances by physical means only by a chemical process The breakdown of a pure substance results in formation of new substances (i.e. chemical change) For a pure substance there is nothing to separate (its only 1 substance to begin with) Mixtures can be separated by physical means (and also by chemical methods, as well) There are 2 general methods of separation Physical: separation based on physical properties Filtration Distillation Centrifugation Chemical: separation based on chemical properties

Chapter 2: Measurements & Calculations Lecture Notes

Types of Observations Qualitative Descriptive/subjective in nature
Detail qualities such as color, taste, etc. Example: “It is really warm outside today” Quantitative Described by a number and a unit (an accepted reference scale) Also known as measurements Notes on Measurements: Described with a value (number) & a unit (reference scale) Both the value and unit are of equal importance!! The value indicates a measurement’s size (based on its unit) The unit indicates a measurement’s relationship to other physical quantities Example: “The temperature is 85oF outside today”

Application of Scientific Notation
Writing numbers in Scientific Notation Locate the Decimal Point Move the decimal point to the right of the non-zero digit in the largest place The new number is now between 1 and 10 Multiply the new number by 10n where n is the number of places you moved the decimal point Determine the sign on the exponent, n If the decimal point was moved left, n is + If the decimal point was moved right, n is – If the decimal point was not moved, n is 0 Writing Scientific Notation numbers in Conventional form Determine the sign of n of 10n If n is + the decimal point will move to the right If n is – the decimal point will move to the left Determine the value of the exponent of 10 Tells the number of places to move the decimal point Move the decimal point and rewrite the number 3

Measurement Systems There are 3 standard unit systems we will focus on: 1. United States Customary System (USCS) formerly the British system of measurement Used in US, Albania, and a couple other countries Base units are defined but seem arbitrary (e.g. there are 12 inches in 1 foot) 2. Metric Used by most countries Developed in France during Napoleon’s reign Units are related by powers of 10 (e.g. there are 1000 meters in 1 kilometer) 3. SI (L’Systeme Internationale) a sub-set set of metric units Used by scientists and most science textbooks Not always the most practical unit system for lab work

Measurements & the Metric System
All units in the metric system are related to the fundamental unit by a power of 10 The power of 10 is indicated by a prefix The prefixes are always the same, regardless of the fundamental unit When a measurement has a specific metric unit (i.e. 25 cm) it can be expressed using different metric units without changing its meaning Example: 25 cm is the same as 0.25 m or even 250 mm The choice of measurement unit is somewhat arbitrary, what is important is the observation it represents 6

Measurement, Uncertainty & Significant Figures
A measurement always has some amount of uncertainty Uncertainty comes from limitations of the techniques used for comparison To understand how reliable a measurement is, we need to understand the limitations of the measurement To indicate the uncertainty of a single measurement scientists use a system called significant figures The last digit written in a measurement is the number that is considered to be uncertain Unless stated otherwise, the uncertainty in the last digit is ±1 Examples: The measurement: 25.2 cm uncertainty: 0.1 cm The measurement: cm uncertainty: 0.01 cm The measurement: cm uncertainty: cm 10

Rules for Counting Significant Figures
Nonzero integers are always significant Zeros Leading zeros never count as significant figures Captive zeros are always significant Trailing zeros are significant if the number has a decimal point Exact numbers have an unlimited number of significant figures Rules for Rounding Off If the digit to be removed is less than 5, the preceding digit stays the same equal to or greater than 5, the preceding digit is increased by 1 In a series of calculations, carry the extra digits to the final result and then round off Don’t forget to add place-holding zeros if necessary to keep value the same!! 12

Exact Numbers Exact Numbers are numbers that are assumed to have unlimited number of significant figures are considered to be known with “absolute” certainty. You do not need to consider or count significant figures for exact numbers. The following are considered exact numbers for CH100: Counting numbers, such as: The number of sides on a square The number of apples on a desktop Defined numbers such as those used for conversion factors, such as: 100 cm = 1 m, 12 in = 1 ft, 1 in = 2.54 cm 1 kg = 1000 g, 1 LB = 16 oz 1000 mL = 1 L; 1 gal = 4 qts. 1 minute = 60 seconds Numbers or constants defined in equations, such as: y = 3x + 15 (both the “3” and the “15” are exact numbers) 14

Converting between Unit Systems
Converting units from one unit system to another (especially within the Metric system) can appear daunting at first glance. However, with a little guidance, and a lot of practice, you can develop the necessary skill set to master this process. To begin, here is a simple mnemonic to guide you through the unit conversion process: Eliminate Replace Relate All unit conversions, regardless of how complex they appear, involve these 3 simple steps. In the following sections, you will be stepped through the unit conversion process using these 3 words as a guide.

Example: Unit Conversion
Convert 25.0 m to cm Convert 1.26 g to kg

Metric Prefixes

Mass Mass is the quantity of matter in a substance
Mass is measured in units of grams Mass does not reflect how much volume something has The kilogram (kg) unit is the preferred unit of mass in the SI system. 1 kilogram is equal to the mass of a platinum-iridium cylinder kept in a vault at Sevres, France. 1 kg has the weight equivalent (on Earth) of lb Conservation of Mass: The total quantity of mass is never created nor destroyed during a chemical process

Distinguishing Mass vs. Weight
The terms mass and weight are commonly used interchangeably but they are fundamentally different! The following are some important differences between mass and weight: Mass is a fundamental property of matter, it is the amount of “stuff” in an object Mass represents an object’s inertia (tendency to resist change in motion) Mass is the same everywhere in the universe SI Units of mass are kilograms (kg) Weight is the effect (or force) of gravity on an object’s mass Weight depends on location (& local gravity) Weight is not a fundamental property of matter SI units of weight are newtons (N) USCS units are pounds (lb)

Volume area height height width length
Volume is the 3-dimensional space that an object occupies Volume Units: The SI unit for volume is the cubic meter, or m3 (meters x meters x meters) The more common metric unit of volume is the Liter (L) In the laboratory, the milliliter (mL) is often more convenient Note: mass and volume are not the same thing (try not to confuse them…). Two objects with the same volume (e.g. a pillow & a sack of potatoes can have different masses and vice versa) length width height height area

Density Density is a property of matter representing the mass per unit volume For equal volumes, a denser object has greater mass For equal masses, a denser object has smaller volume Commonly used units: Solids = g/cm3 (Note: 1 cm3 = 1 mL) Liquids = g/mL Gases = g/L Useful Notes on Density: Volume of a solid can be determined by water displacement Density of matter in various states: solids > liquids >>> gases (exception: water) In a heterogeneous mixture, the denser matter will tend to sink to the bottom 23

Manipulating the Density Equation
mass density volume 24

Chapter 3: Elements & Compounds Lecture Notes

Chemical Symbols & Formulas
Each element has a unique chemical symbol Examples of chemical symbols: Hydrogen: H Oxygen: O Aluminum: Al Each molecule has a unique chemical formula The chemical formula of a molecule indicates the chemical symbol for each of the elements present The # of atoms of each element present in the molecule Examples of chemical formulas: Elemental oxygen: O2 (2 O atoms per molecule) Water: H2O (2 H atoms & 1 O atom) Aluminum sulfate: Al2(SO4)3 (2 Al, 3 S & 12 O atoms)

The Periodic Table All of the known elements are arranged in a chart called the Periodic Table Each element in the Periodic Table is identified by both its chemical symbol and its Atomic Number The elements are organized left-to-right and top-to-bottom according to their Atomic Number The elements are arranged by similarity of chemical properties The columns are called Groups Elements of each group typically have similar properties The rows are called Periods, and reflect the periodicity of chemical properties as atomic number increases

The Periodic Table of Elements

Elements and the Periodic Table
The elements can be categorized as: Metals The leftmost elements in the periodic table Roughly 70% of all of the elements are metals Nonmetals The rightmost elements of the periodic table Semimetals (metalloids) The elements that reside along the “stair step” between the metals and nonmetals in the Periodic Table The properties of semimetals are not quite metallic or non-metallic, but rather somewhere in between

Dmitri Mendeleev (1834-1907) Russian born chemist
Considered one of the greatest science teachers of his era He organized the known elements of his time into the first “periodic table” Elements were organized by chemical properties (& by weight) -> called periodic properties Surprisingly, his periodic table predicted the existence of 3 new elements (which were subsequently discovered)

Chapter 4: Properties of Matter Lecture Notes

Physical & Chemical Properties
Physical Properties are the characteristics of matter that can be changed without changing its composition These characteristics are directly observable or measurable Types of Physical Properties: Extrinsic Physical Properties are unique to objects (i.e. size, shape, mass, etc.) Intrinsic Physical Properties are unique to substances (i.e. density, conductivity, color, etc.) Chemical Properties are the characteristics of a substance that determine the tendency of the matter to transform in composition as a result of the interaction with other substances, the influence of energy or both These are characteristics that describe the behavior of matter 4

Physical & Chemical Changes
Physical Changes are changes that do not result in a change the fundamental composition of the substance Typical Examples: Physical State Changes: boiling, melting, condensing, etc. Shape, Size or Texture Changes Chemical Changes involve a change in the fundamental composition of the matter Notes on Chemical Change: Production of a new substance(s) Referred to as chemical reactions The basic representation: Reactants  Products Note: Both physical and chemical changes will likely produce an alteration of appearance, the key is to discern the type of change that has occurred 8

Energy Energy is loosely described as the capacity of something to do work (or alter the physical or chemical state of an object or system) Common Forms of Energy mechanical, chemical, thermal, electrical, radiant, nuclear The SI unit of energy is the Joule (J) Other commonly used units are Calories (cal) and Kilowatt-hours (kW.hr) Types of energy: Potential: stored energy Kinetic: energy associated with motion and vibration Heat: energy that flows from high to low temperature Principle of Energy Conservation: energy is never created nor destroyed (but it does change from one type to another!)

Distinguishing Heat Energy & Temperature
Temperature is _____. How hot or cold something is (an extrinsic physical property), it represents a particular thermal state Related to the average (kinetic) energy of the substance (not the total energy but the average energy) Measured in units of: Degrees Fahrenheit (oF) Degrees Celsius (oC) Kelvin (K) Heat is _____. Energy that flows from hot objects to cold objects. Heat is not a physical property. Energy absorbed or released by an object resulting in its temperature change Joules (J) Calories (Cal) Kilowatt Hours (kW.hr) Bottom Line: Heat energy absorbed or released is measured by changes in temperature but do not confuse heat energy for temperature

Temperature Scales The 2 traditional temperature scales, Fahrenheit and Celsius, were originally defined in terms of the physical states of water at sea level: Fahrenheit Scale, °F For water: freezing point = 32°F, boiling point = 212°F Celsius Scale, °C For water: freezing point = 0°C, boiling point = 100°C 1 Celsius temperature unit is larger than 1 Fahrenheit unit The SI unit for temperature is a variant of the Celsius scale Kelvin Scale, K For water: freezing point = 273 K, boiling point = 373 K The Kelvin temperature unit is the same size as the Celsius unit 22

Temperature of ice water and boiling water.

Heat (Energy) Heat is energy that flows due to a temperature difference Heat energy flows from higher temperature to lower temperature Heat is transferred due to “collisions” between atoms/molecules of different kinetic energy When produced by friction, heat is mechanical energy that is irretrievably removed from a system Processes involving Heat: Exothermic = A process that releases heat energy. Example: burning paper is an exothermic process because energy is produced as heat (the temperature rises!). Endothermic = A process that absorbs energy. Example: melting ice to form liquid water is an endothermic process because heat energy must be absorbed to change the physical state (in this case the temperature does not change!). 17

Heat (cont.) When something absorbs or loses heat energy, 1 of 2 things can occur: Its temperature will change (e.g. hot coffee will cool down) Its physical state will change (e.g. ice will melt) For the former case above, the heat energy absorbed or lost by an object is proportional to: The mass of the object (m) The change in temperature the object undergoes (DT) The specific heat capacity (s) (a physical property unique to the substance) To calculate heat gained (Q): Q m DT s

Specific Heat Capacity (s)
Specific heat capacity reflects how absorbed heat energy relates to the corresponding increase in temperature for a given amount of mass, i.e. energy per unit mass per unit temperature change or Specific Heat Capacity is commonly measured in units of: J/goC (SI) cal/goC (metric & more useful in the lab) Specific Heat Capacity is a unique intrinsic physical property of matter. Typically, __________. Metals have low specific heat capacity Non-metals have higher specific heat capacity than metals Water has an unusually large specific heat capacity

0.900 0.473

Table of Specific Heat for Various Substances
J/g.K cal/g.K J/mol.K Aluminum 0.900 0.215 24.3 Iron 0.473 0.113 26.4 Copper 0.385 0.0921 24.5 Brass 0.380 0.092 ... Gold 0.131 0.0312 25.6 Lead 0.128 0.0305 Silver 0.233 0.0558 24.9 Tungsten 0.134 0.0321 24.8 Zinc 0.387 0.0925 25.2 Mercury 0.140 0.033 28.3 Alcohol (ethyl) 2.138 0.511 111 Water 4.184 1.000 75.2 Ice (-10 C) 2.059 0.492 36.9 Granite .790 0.19 Glass .84 0.20

Chapter 5: Early Atomic Theory & Structure Lecture Notes

Early Model of Matter: Aristotle (384-322 BC)
Introduced observation as an important step in understanding the natural world According to his model of nature, all forms of matter are mixtures of one of 4 basic “elements”: All matter has one or more of 4 basic “qualities”: According to Aristotle: Any substance could be transformed into any other substance by altering the relative proportion of these elements and qualities (i.e. lead to gold) 4) Fire 2) Water 3) Air 1) Earth 4) Dry 2) Moist 3) Hot 1) Cold

Dalton’s Atomic Theory
Each element consists of individual particles called atoms Atoms can neither be created nor destroyed All atoms of a given element are identical Atoms combined chemically in definite whole-number ratios to form compounds Atoms of different elements have different masses

The Modern Atomic Model
According to our modern model of the matter, the atom has 2 primary regions of interest: Nucleus Contains protons & neutrons (called nucleons, collectively) Establishes most of the atom’s mass Mass of 1 neutron = x10-27 kg Mass of 1 proton = x10-27 kg Small, dense region at the center of the atom The radius of the nucleus ~ m (1 femtometer) The Electron Cloud Contains electrons Mass of 1 electron = x10-31 kg Establishes the effective volume of the atom The radius of the electron cloud ~ m (1 Angstrom) Determines the chemical properties of the atom During chemical processes, interactions occur between the outermost electrons of each atom The electron properties of the atom will define the type(s) of interaction that will take place

Structure of the Atom

What holds the atom together?
Electromagnetic interaction (a.k.a. electric force) holds the electrons to the nucleus The negative charge (-) of the electrons are attracted to the positive charge (+) of the nucleus Strong interaction (a.k.a. strong force) holds the nucleons together within the nucleus The positive charge of the protons repel each other All nucleons, protons and neutrons, possess a STRONG attraction to each other that overcomes the protons’ mutual repulsion

Electric Charge Electric charge is a fundamental property of matter
We don’t really know what electric charge is but we do know that there are 2 kinds: Positive charge (+) Negative charge (-) Opposite charge polarity is attractive: + attracts - Same charge polarity is repulsive: + repels + and – repels – The magnitude of electric charge (q) is the same for protons and electrons: The charge of a proton (qproton) or electron (qelectron) is the smallest amount that occurs in nature, it is called the quantum of charge: qproton = x Coulombs (or 1+) qelectron = x Coulombs (or 1-)

Ions Atoms (or molecules) that have gained or lost one or more electrons Ions that have lost electrons are called cations Ions that have gained extra electrons are called anions Ionic compounds have both cations and anions (so that their net charge is zero)

Ions (cont.) Ions are electrically charged atoms and thus carry electric charge: The electric charge of an ion is due to the imbalance of electrons and protons When an atom has lost one or more of its electrons it carries a positive charge “1+” for each electron that is lost When an atom has gained one or more of its electrons it carries a positive charge “1-” for each excess electron that is gained When an atom/molecule is an ion, its charge must be specified: Sodium ion: Na+ Chloride ion: Cl- Hydroxide ion: OH- Notes on Electric Charge: Opposite charges attract Like charges repel + - + -

# of neutrons = (Mass #) – (Atomic #)
Atomic Bookkeeping Atomic number (Z) The number of protons in an atom or ion The number that defines the identity of the atom Mass number (A) The number of protons & neutrons in a specific atom or isotope The number that represents the mass of an atom To determine number of neutrons in an atom: # of neutrons = (Mass #) – (Atomic #) Or # of neutrons = A - Z

Mass # vs. Atomic Mass (Atomic Symbol) C C
Isotopes are the equivalent of sibling members of an element Unique atoms of the same element with different mass numbers (i.e. they have different numbers of neutrons) Unique isotopes are identified by their mass number Isotope notation: Example: carbon-12 ( ) & carbon-14 ( ) Atomic mass The average total mass of an element’s various naturally occurring isotopes The unit of Atomic Mass is the Dalton (or amu) 1 Dalton = one twelfth mass of one 12C atom = 1.661x10-27 kg Note: There 6 protons & 6 neutrons in a 12C atom but the mass of a 12C atom is actually slightly less than the combined mass of all of the nucleons individually. Where is this lost mass? It’s released as energy when the nucleons combine (bind) to form the nucleus of the atom. (Atomic Symbol) Mass # Atomic # C 12 6 C 14 6

Examples of Isotopes

Chapter 6: Nomenclature of Inorganic Compounds

Types of Compounds When compounds are formed they are held together by the association of electrons This association is called a chemical bond There are 3 general types of chemical bonds: Ionic Covalent (or molecular) Polar covalent Simple compounds are classified (and thus named) according to the type of chemical bond(s) that hold together its atoms Note: many compounds have more than one type of chemical bond present, but we will focus on only “simple compounds”

Types of Compounds (cont.)
For “practical” purposes will separate all simple compounds into 2 general categories: Ionic Compounds Made up of ions (both positive and negative charge) Must have no net charge (i.e. combined charge of zero) Depend on the attraction between positive and negative charges of the ions Usually a metal is present as a cation and a nonmetal is present as an anion Non-Ionic (aka: Molecular or Covalent) Compounds Made up of atoms that share their outer electrons Electric charge plays no direct role in their formation There are usually no metals are present in these compounds

Naming Compounds The easiest way (usually) to identify an ionic compound is to ask whether or not there is a metal present in the chemical formula (or the name): Is a metal present? Yes -> it is an Ionic Compound (e.g. CaCl2) No -> it is a Non-Ionic Compound (e.g. CCl4) or an Acid Notes: Ionic compounds do not use the Greek prefixes and are named according to the identity of the ions present Non-Ionic compounds require the use of Greek prefixes to indicate the number of each element present in one molecule

Naming Simple Compounds
A “simple” or binary compound is a compound made of only 2 types of elements When the first element is a metal: The first element (metal) keeps its full name The non-metal goes by its root with the suffix “-ide” added to the end Example: NaCl is sodium chloride When there are no metals present Same as above except Greek prefixes must be used to identify the number of each element present in the compound Example: CO2 is carbon dioxide

Determining Chemical Formula of an Ionic Compound
To determine the chemical formula of an ionic compound from its chemical name: Identify the ions present, both cation(s) and anion(s), from the name. Example: potassium sulfide Cation: potassium Anion: sulfide Determine the ionic charge of the ions Example: {from above} potassium ion, K+ sulfide ion, S2- Determine the number of each ion needed to obtain a neutral compound Example: {from above}  2 K+ ions are needed for every S2- Combine the chemical sysmbols of the ions to get the final chemical formula Example: {from above}  K2S is the formula for potassium sulfide

Ionic Charges & the Periodic Table
The position of an element in the Periodic Table is a useful indicator of the type of ion an element is capable of forming: Group 1 metals form 1+ cations (Na+ sodium ion) Group 2 metals form 2+ cations (Ca2+ calcium ion) Group 13 metals form 3+ cations (Al3+ aluminum ion) Group 3-12 Metals (plus Sn, Pb, & Bi) can form more than one type of cation Roman numerals are used to indicate the charge of the cation Example: Fe3+ is called iron(III) FeCl3 is called iron(III) chloride Notable Exceptions: Ag+, Cd2+ & Zn2+ Group 15 nonmetals form 3- anions (e.g. N3- nitride ion) Group 16 nonmetals form 2- anions (e.g. O2- oxide ion) Group 17 nonmetals form 1- anions (e.g. Cl- chloride ion) Group 18 elements do not form ions

Greek Prefixes for Compound Names
1) Mono- 2) Di- 3) Tri- 4) Tetra- 5) Penta- CCl4 is carbon tetrachloride 6) Hexa- 7) Hepta- 8) Octa- 9) Nona- 10) Deca- C3H8 is tricarbon octahydride Notes: 1) Prefixes are used when the compound does not have a metal present (or when H is the first element in the formula) 2) Prefixes must be used for every element present in the compound 3) Mono- is not used for the first element in a compound name (e.g. carbon dioxide)

Ionic Compounds containing Polyatomic ions
Some ionic compounds are made up of polyatomic ions Polyatomic ions are usually ions formed from non-ionic molecules e.g. The sulfate ion, SO42-, is essentially a molecular compound containing S and O with 2 additional electrons When you encounter polyatomic ions in compounds, do not freak out!! Become familiar with the common polyatomic ions on the handout Example: The nitrate ion (NO3-) Fortunately, the naming of ionic compounds containing polyatomic ions is similar to that for ionic compounds

KOH(aq) + H+  K+(aq) + H2O (l)
Acids From the Latin term for “sour” {Acids are sour to the taste} Acids are substances that donate or release hydrogen cations, H+, (usually when dissolved in water) The chemical formula for acids usually begins with H Example: hydrochloric acid (HCl) HCl(aq)  H+ + Cl- (aq) Taste bitter (Note: it is not advised to taste strong bases…) Usually metal containing hydroxides Substances that accept hydrogen cations (H+) when dissolved in water Example: potassium hydroxide (KOH) KOH(aq) + H+  K+(aq) + H2O (l) Bases

Naming Acids Lets separate acids into 2 types:
Acids that contain oxygen Acids that do not contain oxygen Naming acids containing oxygen: For acids containing “-ate” anions: Use root of the anion (for sulfate, SO42-, use sulfur) Add “-ic” suffix then end with “acid” Example: H2SO4 is sulfuric acid For acids with “-ite” anions: Use root of the anion (for sulfite, SO32-, use sulfur) Add “-ous” suffix then end with “acid” Example: H2SO3 is sulfurous acid

Naming Acids (cont.) Naming acids not containing oxygen:
Add “hydro-” prefix to beginning Use root of the anion (i.e. Cl- use chlor) Add “-ic” suffix then end with “acid” Example: HCl is hydrochloric acid Name the following acids: HF HNO2 HCN H3PO4

Antoine Lavoisier ( ) Considered by many to be the “Father of Modern Chemistry” Major contributions included Demonstrated that water cannot be transmuted to earth Established the Law of Conservation of Mass Developed a method of producing better gunpowder Observed that oxygen and hydrogen combined to produce water (dew) Invented a system of chemical nomenclature (still used in part today!) Wrote the 1st modern chemical textbook

Ch 7: Quantitative Composition of Compounds Lecture Notes (Sections 7.1 to 7.3)

The Mole The mole is a counting unit (analogous to the dozen unit)
A large unit used to describe large quantities such as number of atoms 1 mole = x 1023 units 6.022 x 1023 is known as Avogadro’s number (NA) Relationship between the mole & the Periodic Table The atomic mass is the quantity (in grams) of 1 mole of that element The units of atomic mass are grams/mole Mass is used by chemists as a way of “counting” number of atoms/molecules of a substance Mole calculations

What do you get if you have Avogadro's number of donkeys?
Got mole problems? Call Avogadro at What do you get if you have Avogadro's number of donkeys? Answer: molasses (a mole of asses)

Molar Mass Molar mass is the mass in grams of 1 mole of a substance
Molar mass refers to both atoms & molecules Elements (atoms) Examples: 1 mole of Na has a mass of g 1 mole of Cl has a mass of 35.45 1 mole of Cl2 has a mass of g Compounds (molecules) 1 mole of NaCl has a mass of g Mass of Na (22.99 g) + Mass of Cl (35.45 g) 1 mole of CO2 has a mass of g Mass of C (12.01 g) + 2 x Mass of O (16.00 g)

Mole Calculations To convert from atoms (or molecules) to moles, divide the # of atoms (or molecules) by Avogadro’s # Example: How many moles are 1.0x1024 atoms? To convert from moles to atoms (or molecules), multiply the # of atoms (or molecules) by Avogadro’s # Example: How many molecules are in 2.5 moles?

Mole-Mass Calculations
To convert from moles to grams, multiply the # of moles by atomic mass Example: How many grams in 2.5 moles of carbon? To convert from grams to moles, divide the mass in grams by atomic mass Example: How many moles are in 2.5 g of lithium?

Percent Composition Percent composition is the percentage of each element in a compound (by mass) Percent composition can be determined from either: the formula of the compound the experimental mass analysis of the compound Note: The percentages may not always total to 100% due to rounding 12

Percent Composition Calculations
To determine % Composition from the chemical formula: Determine the molar mass of compound Multiply the molar mass of the element of interest by the number of atoms per molecule then Divide this value by the molar mass of the compound Example: The % Composition of sodium in table salt The molar mass of NaCl is g/mol There is 1 atom of Na in each NaCl molecule The atomic mass of Na is 22.99

Percent Composition Calculations
Perform the following % Composition calculations: The % composition of carbon in carbon monoxide The % composition of oxygen in water The % composition of chlorine in sodium hypochlorite

Amadeo Avogadro (1743-1794) Italian lawyer turned chemist
Major contributions included: Established difference between atoms & molecules: Oxygen & nitrogen exist as molecules O2 & N2 Reconciled the work of Dalton & Guy-Lussac Establishing Avogadro’s Principle: equal volumes of all gases at the same temperature and pressure contain the same number of molecules. Note: Avogadro did not determine Avogadro’s number nor the mole (these concepts came later) Avogadro is honored because the molar volume of all gases should be the same Much of Avogadro’s work was acknowledged after he died, by Stanislao Cannizarro

Chapter 8: Chemical Equations Lecture Notes

Chemical Equations (Intro)
Chemical equations are used to symbolically describe chemical reactions In a chemical equation (or reaction for that matter) the substances that undergo chemical change(s) are called the reactants The resulting substances formed are called the products The standard representation of a chemical equation: Reactant(s)  Product(s) Example: The production of water 2H2 (g) + 1O2 (g)  2H2O (g) The underlined numbers are called coefficients. The number of each molecule for each reactant & product in the chemical reaction They are always whole numbers

Chemical Equations (cont.)
Balanced chemical equations indicate the ____ identity of each reactant & product involved in the reaction phase of each reactant and product involved in the reaction (i.e. solid (s), liquid (l) or gas (g)) relative quantity of each reactant and product involved in the reaction (the coefficients!) relative molar quantity of each reactant and product involved in the reaction (the coefficients!)

Balancing Chemical Equations
According to the Law of Mass Conservation (& John Dalton!) matter is never created nor destroyed during chemical reactions All of the atoms in the reactants of a chemical reaction must be accounted for in the products The Basic Process of Balancing Chemical Equations: Identify all reactants & products in the reaction & write out their formulas (this is the unbalanced chemical equation) Count the number of each atom for each compound for each reactant & product (these values must be the same for both reactants & products when the reaction is balanced!) Starting with the most “complicated” molecule, systematically adjust the coefficients to balance # of the atoms on each side of the reaction (balance one atom at a time) Repeat until all atoms are balanced for the reaction Now you should have a balanced chemical equation!

Balancing Chemical Equations (example)
When sodium metal is added to water a violent reaction takes place producing aqueous sodium hydroxide and releasing hydrogen gas. Write out the unbalanced chemical reaction: Now, balance the chemical reaction:

Balancing Chemical Reactions (Hint)
When a polyatomic ion(s) appears on both the reactant & product side of the reaction unchanged, treat the whole ion as a “unit” when balancing the reaction Example: Note the nitrate ion (NO3-) gets swapped between the Ag+ and the Ca2+ ions in this reaction So NO3- can be treated as a whole unit when balancing this reaction Balance it! AgNO3(aq) + CaCl2(aq) AgCl(s) + Ca(NO3)2(aq)

Common Classifications for Chemical Reactions
Combination (or Synthesis): reactions in which reactants combine to make one product Decomposition: reactions in which one reactant breaks down into smaller products Single Displacement: reactions where a part of one reactant is displaced and combined with another reactant 2Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) Double Displacement: reactions where a part of two reactants is displaced and exchanged AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) Examples: Acid-base neutralization Formation of insoluble products (Precipitation reactions) Metal oxide + acid Gas formation Oxidation-Reduction Reactions: reactions involving the transfer or rearrangement of electrons 18

Combination & Decomposition Reactions
Reactions in which chemicals combine to make one product are called Combination or Synthesis Reactions Metal + Nonmetal reactions can be classified as Combination Reactions 2 Na(s) + Cl2(g)  2 NaCl(s) Reactions between Metals or Nonmetals with O2 can be classified as Combination Reactions N2(g) + O2(g)  2 NO(g) Note: these two types of Combination Reactions are also subclasses of Oxidation-Reduction Reactions Reactions in which one reactant breaks down into smaller molecules are called Decomposition Reactions Decomposition reactions are generally initiated by the addition of energy (via electric current or heat) Decomposition reactions are the opposite of Combination Reactions: 2 NaCl(l)  2 Na(l) + Cl2(g) 23

Single Displacement Reactions
Single displacement reactions involve one part of a reactant being transferred to another The basic pattern of the single displacement reaction: XY + A X + AY Example 1: Metal + Acid  Salt + Hydrogen Zn(s) + 2 HCl(aq)  ZnCl2(aq) + H2(g) Example 2: Metal + Water  Hydrogen + Metal Oxide (or metal hydroxide) 3 Fe(s) + 4 H2O(l)  4 H2(g) + Fe2O3(s) Example 3: Metal + Salt  Metal + Salt 2 Al(s) + Fe2O3(s)  2 Fe(s) + Al2O3(s) Example 4: Halogen + Halide Salt  Halogen + Halide Salt Cl2(g) + 2 NaBr(s)  Br2(g) + 2 NaCl(s) 21

Double Displacement Reactions
Double Displacement Reactions involve the double exchange of a component (such as ions) between two reactants The basic form of double displacement reactions is: XY + AB  XB + AY where X, Y, A, and B are the components of the reactants Example 1: Acid Base Neutralization H2SO4(aq) + Ca(OH)2(aq)  CaSO4(aq) + 2 H2O(l) Example 2: Metal Oxide + Acid CaO(s) + 2HCl(aq)  CaCl2(aq) + H2O(l) Example 3: Formation of an Insoluble Precipitate (Precipitation) KCl(aq) + AgNO3(aq)  KNO3(aq) + AgCl(s) Example 4: Formation of a Gas HCl (aq) + ZnS (s)  ZnCl2(aq) + H2S (g) or 2HOH or HOH 19

Solubility & Precipitation Reactions
When 2 solutions are combined and result in the formation of an insoluble product: The product will not dissolve in the solvent The product will form a precipitate Solubility is an intrinsic physical property and a measure of how well a substance (solute) will dissolve in another substance (solvent) Solubility is temperature dependent Solid solubility increases with increased temperature (i.e. you can dissolve more sugar in hot water than in cold water) Gas solubility increases with decreased temperature (i.e. you can dissolve more CO2 in cold water than hot water) A solute is soluble if any of it will dissolve in a solvent Eg. NaCl is soluble in water A solute is insoluble if no appreciable amount of it will dissolve in solvent Eg. AgCl is insoluble in water Precipitation (formation of an insoluble solid) is one indication that a chemical change has occurred!

General Rules for Solubility
Most compounds that contain NO3- ions are soluble Most compounds that contain Na+, K+, or NH4+ ions are soluble Most compounds that contain Cl- ions are soluble, except AgCl, PbCl2, and Hg2Cl2 Most compounds that contain SO42- ions are soluble, except BaSO4, PbSO4, CaSO4 Most compounds that contain OH- ions are slightly soluble (will precipitate), except NaOH, KOH, are soluble and Ba(OH)2, Ca(OH)2 are moderately soluble Most compounds that contain S2-, CO32-, or PO43- ions are slightly soluble (will precipitate) 8

Oxidation-Reduction Reactions
Reactions that involve transfer or rearrangement of electrons are called oxidation-reduction reactions. Examples of oxidation-reduction reactions: Metal + Nonmetal: 2Na(s) + Cl2(g)  2NaCl(s) The metal loses an electron(s) and becomes a cation (oxidation  metal gets oxidized: Na  Na+ + e-) The nonmetal gains an electron(s) and becomes an anion (reduction  nonmetal gets reduced: Cl + e-  Cl-) In this reaction, electrons are transferred from the metal to the nonmetal O2 as a reactant or product: CH4(s) + O2(g)CO2(g) + H2O(g) In this reaction, it is not obvious that electron transfer has taken place. In this case, oxidation states are altered. Often this type of reaction involves the release of large amounts of energy, even combustion 20

Rates of Chemical Reactions
How quickly a chemical reaction occurs is indicated by its reaction rate How quickly the concentration of products increases How quickly the concentration of reactants decreases The Factors that influence reaction rates: Reactants must be in contact Reactions occur due to collisions Without contact between reactants there can be no reaction Concentration of reactants The more reactant molecules packed into a given space the more likely a collision (& reaction) will occur Temperature the average KE of each reactant affects how much energy will be transferred between reactants during a molecular collision Molecules must transfer enough KE to break the existing bonds Energy due to collisions must be great enough to overcome the energy barrier for existing chemical bonds. The magnitude of this energy is referred to as activation energy.

The Role of Energy in Chemical Reactions
Energy transformations always accompany chemical reactions: Energy is required to break bonds (energy absorbed or activation energy) Energy is released when bonds are formed Note: The amount of energy required to break a chemical bond is the same as the energy released when that type of bond is formed, this is called the Bond Energy For a chemical reaction to occur: Energy must be absorbed in order to break chemical bonds in the reactants Energy is released as new bonds are formed in the products Endothermic reactions absorb more energy than they release N2(g) + O2(g) kJ  2NO(g) Exothermic reactions release more energy than they absorb H2(g) + Cl2(g)  2NO(g) kJ

Energy in Chemical Reactions
Potential Energy Reactants Products Activation Energy (EA) Energy Released (Q) Exothermic Reactions Potential Energy Reactants Products Activation Energy (EA) Energy Absorbed (Q) Endothermic Reactions

Energy in Reactions (cont.)
Internal Energy 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g) Low Activation Energy (EA) Large amount of Energy Released (Q) Example: Sodium Water Reaction

Catalysts Catalysts are substances that speed up chemical reactions
Allow reactions to occur that might not otherwise take place (due to low temperature for example) Lower activation energy for a chemical reaction Participation of catalysts in a chemical reaction They may undergo a chemical change as a reactant but they are always recycled as a product (so there is no net change in the catalyst molecule) Catalysts are indicated in a chemical reaction by placing the chemical formula over/under the reaction arrow. Reactants  Products Example: The breakdown of hydrogen peroxide catalyst

Catalysts & Energy in Reactions
Potential Energy Reactants Products Activation Energy with catalyst Catalysts lower Activation Energy Activation Energy without catalyst

Ch 9: Calculations from Chemical Reactions Lecture Notes (Sections 9.1 to 9.5)

Chemical Equations: What do they tell us?
A properly written chemical equation will provide the following information: All reactants & products involved in the reaction The physical state of all reactants & products The presence of any catalysts involved in the chemical reaction The relative quantity of all reactants & products Molecule to molecule ratios Mole to mole ratios Even mass to mass ratios can be determined (with use of molar mass values)

Information Given by the Chemical Equation
A balanced chemical equation provides the relationship between the relative numbers of reacting molecules and product molecules Example: The formation of carbon dioxide from carbon monoxide and oxygen gas 2 CO + O2  2 CO2 In this chemical equation, it is indicated that 2 CO molecules react with every 1 O2 molecule to produce 2 CO2 molecules Alternative interpretation: there is a 2:1 (numerical) ratio of CO to O2 for this completed reaction, 2 CO:1 O2 :2 CO2 2

Interpretation of the Chemical Equation
Since the information given in a balanced chemical reaction is relative: 2 CO + O2  2 CO2 the following are alternative interpretations of the chemical equation: 200 CO molecules react with 100 O2 molecules to produce 200 CO2 molecules 2 billion CO molecules react with 1 billion O2 molecules to produce 20 billion CO2 molecules 2 moles CO molecules react with 1 mole O2 molecules to produce 2 moles CO2 molecules 12 moles CO molecules react with 6 moles O2 molecules to produce 12 moles CO2 molecules Note: The coefficients in the balanced chemical equation also shows the molecules and mole ratio of the reactants and products Since moles can be converted to masses, we can determine the mass ratio of the reactants and products as well 3

Mole and Mass Ratios in Chemical Equations
For the following chemical equation: 2 CO + O2  2 CO2 The following mole relations are implied: 2 moles CO : 1 mole O2 : 2 moles CO2 Note the molar masses of the compounds in this reaction: 1 mole of CO = g 1 mole O2 = g 1 mole CO2 = g The mass ratio of the compounds in this reaction can be determined using the molar mass values: 2(28.01) g CO : 1(32.00) g O2 : 2(44.01) g CO2 The mass ratio of the compounds in this reaction are: 56.02 g CO : g O2 : g CO2 5

2 moles CO : 1 mol O2 : 2 moles CO2
Example Determine the Number of Moles of Carbon Monoxide required to react with 3.2 moles Oxygen, and the moles of Carbon Dioxide produced Write the balanced equation 2 CO + O2  2 CO2 Use the coefficients to find the mole relationship 2 moles CO : 1 mol O2 : 2 moles CO2 Use dimensional analysis to obtain the # of moles The # mol of CO: The # mol of CO2: 6

Ch 100: Fundamentals for Chemistry

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