General Chemistry CHMY 121N - 00
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This 18 page Study Guide was uploaded by Gabby Goodell on Sunday September 18, 2016. The Study Guide belongs to CHMY 121N - 00 at University of Montana taught by Daniel J. Dwyer (P) in Fall 2015. Since its upload, it has received 5 views. For similar materials see Intro to General Chemistry in Chemistry at University of Montana.
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STUDY GUIDE FOR MIDTERM #1 LECTURE 1 Chemistry is the science of matter. Chemistry deals with the composition, structure, properties, reactions and energetics of matter. Matter is any object that has mass and occupies space. Water (H2O) is composed of two hydrogen atoms and one oxygen atom. The atoms must connect in a specific arrangement, in this case forming two O-H bonds. Scientists use a logical process to explain the world around them. This process is called the Scientific Method. Collect the facts and data relevant to your question. 2. Formulate a hypothesis. 3. Plan and perform additional experiments to test the hypothesis. 4. Modify the hypothesis, perform an experiment, then analyze the resulting data to look for trends that relate to the question. A hypothesis is a tentative explanation of the data that requires further experimentation to be validated. A useful hypothesis must explain all of the data. A well-established hypothesis is often called a theory. A theory summarizes a hypothesis that has been supported by repeated experimentation. A theory is valid as long as there is no evidence to disprove it. A scientific law is a statement of natural phenomenon where no exceptions are known under the given conditions. Matter appears in many forms, both big and small. On the microscopic level, all matter is composed of discrete, tiny fundamental particles called atoms. Matter exists as three physical states: solid, liquid and gas. A solid has definite shape and volume which can be independent of its container. o The most common solids are crystalline and have regular, repeating three-dimensional geometric patterns. o Solid water molecules are held together rigidly and are very close to each other. o Some solids such as plastics, glass and gels do not have regular, internal geometric patterns. o These solids are called amorphous solids, meaning without shape or form. A liquid has definite volume but not a definite shape. o Liquid particles are held together by strong attractive forces and are able to move freely. o Liquid water molecules are close together but are free to move around and slide over each other. o Liquids are fluid which allows them to take the shape of the container. A gas has indefinite volume and no fixed shape. o Gas particles move independently and are relatively far apart, which allows them to completely fill a container. o Gaseous water molecules are far apart and move freely and randomly. o Gases can be compressed or expanded almost indefinitely. Although matter is separated into discrete units, attractive forces hold the particles together and give matter its appearance of continuity. o Attractive forces are strongest in solids, weaker in liquids and weakest in gases. The physical state of a substance depends on its temperature and pressure. A pure substance has a definite, fixed composition and is either an element or a compound. o The sugar on the spoon and the water in the beaker are each a pure substance – both are compounds. Homogeneous matter is uniform in appearance and has the same properties throughout. Heterogeneous matter consists of two or more physically distinct phases. o Ice floating in water is a two phase system. Each phase is homogeneous but the overall system is heterogeneous. o A phase is a homogeneous part of a system separated from other parts by a physical boundary. A mixture is a combination of two or more pure substances and can be homo- or heterogeneous. o Sugar dissolved in water is a homogeneous mixture. o The proportion of sugar and water can be varied but the composition will be the same throughout. o A heterogeneous mixture will have a different composition depending on where the sample is taken. The components of a mixture do not lose their identities and may be separated by physical means such as: boiling, filtration, floatation, magnetism Sulfur and iron can be separated using a magnet. LECTURE 2 Precision and accuracy Precision: Reproducibility of a measurement. o Precision depends on how well you can read an instrument and on how stable the measurement conditions are. Accuracy: Measure of closeness to correct value o Accuracy of a measurement depends on how well your instrument is preforming. o The accuracy of most scientific measurements is unknown to increase the probability of accuracy. Scientific Notation: A way to write very large or small numbers (measurements) in a compact form. o Number written from 1-10 Raised to a power (-/+ or fractional) o Move the decimal point in the original number so that it is located after the first nonzero digit. o 2. Multiply this number by 10 raised to the number of places the decimal point was moved. o 3. Exponent sign indicates which direction the decimal was moved. EXAMPLE Write 0.000423 in scientific notation. Place the decimal between the 4 and 2. 4.23 The decimal was moved 4 places so the exponent should be a 4. 4.23 x 10-4 The decimal was moved to the right so the exponent should be negative, if it was moved to the left it would be positive. Because all measurements involve uncertainty, we must be careful to use the correct number of significant figures in calculations. Rules for Counting Significant Figures o All nonzero digits are significant. o 2. Some numbers have an infinite number of sig figs Ex. 12 inches are always in 1 foot Exact numbers have no uncertainty. o 3. Zeroes are significant when: o a. They are in between non zero digits Ex. 75.04 has 4 significant figures (7,5,0 and 4) o b. They are at the end of a number after a decimal point. o Ex. 32.410 has five significant figures (3,2,4,1 and 0) o 4. Zeroes are not significant when: o a. They appear before the first nonzero digit. Ex. 0.00321 has three significant figures (3,2 and 1) o b. They appear at the end of a number without a decimal point. o Ex. 6920 has three significant figures (6,9 and 2) When in doubt if zeroes are significant, use scientific notation! LECTURE 3 Metric or International System (SI): Standard system of measurements for mass, length, time and other physical quantities. Based on standard units that change based on factors of 10. Prefixes are used to indicate multiples of 10. This makes the metric system a decimal system. Meter (m): standard unit of length of the metric system. Common Length Relationships: 1 meter (m) = 100 centimeters (cm) 1 kilometer (km) = 1000 meters Relationship Between the Metric and English System: 1 inch (in.) = 2.54 cm = 1000 millimeters (mm) Conversion factor: A ratio of equivalent quantities. Dimensional analysis: converts one unit of measure to another by using conversion factors. Example: 1 km = 1000 m Conversion factors can always be written two ways. Both ratios are equivalent quantities and will equal 1. Any unit can be converted to another unit by multiplying the quantity by a conversion factor. Unit1 x conversion factor = Unit2 Example A conversion factor must cancel the original unit and leave behind only the new (desired) unit. The original unit must be in the denominator and new unit must be in the numerator to cancel correctly. o 1 km 1000 m 2 m x = 0.002 km Units are treated like numbers and can cancel. A systematic method to solve these types of numerical problems is key. Convert 215 centimeters to meters. Solution Map: known quantity desired quantity 1 m 100 cm = 2.15 m215 cm x Convert 125 meters to kilometers. 1 km 1000 m = 0.125 km125 m Some problems require a series of conversions to get to the desired unit. Each arrow in the solution map corresponds to the use of a conversion factor. How many feet are in 250 centimeters? How many cm3 are in a box that measures 2.20 x 4.00 x 6.00 inches? Mass: amount of matter in an object Mass is measured on a balance. Weight: effect of gravity on an object. Mass is independent of location, but weight is not. Weight is measured on a scale, which measures force against a spring. Mass is the standard measurement of the metric system. The SI unit of mass is the kilogram. (The gram is too small a unit of mass to be the standard unit.) How many centigrams are in 0.12 kilograms? Volume: the amount of space occupied by matter. The SI unit of volume is the cubic meter (m3) The metric volume more typically used is the liter (L) or milliliter (mL). A liter is a cubic decimeter of water (1 kg) at 4 °C. Volume can be measured with several laboratory devices. Common Volume Relationships 1 L = 1000 mL = 1000 cm3 1 mL = 1 cm3 1 L = 1.057 quarts (qt) The composition of many mixtures is often given in percent. Percent can be defined as parts per 100 x parts = x where x equals percent 100 total parts Mass percent Use the formula mass percent = mass part x 100% mass total Since the same units cancel out any mass units can be used in the formula Thermal energy: A form of energy involving the motion of small particles of matter. Temperature: measure of the intensity of thermal energy of a system (i.e. how hot or cold). Temperature is measured using a thermometer. Temperature can be expressed in 3 commonly used scales. o Celsius (°C), o Fahrenheit (°F) o Kelvin (K). Celsius and Fahrenheit are both measured in degrees, but the scales are different. The Fahrenheit scale has a range of 180° between freezing and boiling. H2O °C °F K Freezing Point 0 °C 32 °F 273.15 K Boiling Point 100 °C 212 °F 373.15 K The lowest temperature possible on the Kelvin scale is absolute zero (-273.15 °C). Heat: flow of energy due to a temperature difference. Heat flows from regions of higher to lower temperature. The SI unit of temperature is the Kelvin (K). Density (d): the ratio of the mass of a substance to the volume occupied by that mass. Density is a physical property of a substance. The units of density are generally expressed as g/mL or g/cm3 for solids and liquids and g/L for gases. The volume of a liquid changes as a function of temp, so density must be specified for a given temperature. Ex. The density of H2O at 4 ºC is 1.0 g/mL while the density is 0.97 g/mL at 80 ºC. Specific gravity (sp gr): ratio of the density of a substance to the density of another substance (usually H2O at 4 ºC). Specific gravity is unit-less (in the ratio all units cancel). LECTURE 4 An element is a fundamental substance that cannot be broken down by chemical means into a simpler substance. Elements are the building blocks of matter. Elements can occur naturally or be synthesized in labs. The smallest unit of an element that retains its properties and chemical behavior is called an atom. Currently, 118 elements are known. 88 of these elements are naturally occurring. At room temperature, only bromine (Br) and mercury (Hg) are liquids. At room temperature, 11 elements are gases: hydrogen (H2), nitrogen (N2), oxygen (O2), fluorine (F2), chlorine (Cl2), helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn). Ten elements make up almost 99 % of the mass of the Earth’s crust, seawater and atmosphere. Oxygen accounts for about 20 % of the atmosphere and is found in nearly all rocks, sand and soil. The most prevalent elements in the human body are oxygen (65 %), carbon (18 %) hydrogen (10 %). Many element names are derived from Greek, Latin or German words that describe a property of the element. Iodine comes from the Greek word iodes meaning “violetlike” and iodine is violet colored as a vapor. Some elements are named for the location of their discovery, such as germanium which was discovered in 1886 by a German chemist. Other elements commemorate famous scientists, like curium which is named after Marie Curie. Each element has an abbreviation called a symbol. Fourteen elements have single letter symbols and the rest have two letter symbols. The first letter of a symbol must always be capitalized. If a second letter is needed, it should be lowercase. Consider the element cobalt: Cobalt has the symbol Co Writing CO would symbolize elements C (carbon) and O (oxygen) and would represent carbon monoxide. The Periodic Table was designed by Dimitri Mendeleev in 1869 to organize elements based on their properties. Elements with similar chemical properties are placed in columns called groups. Four groups have special identifying names, o like Noble Gases, which are all unreactive gases. The eight tall columns are called representative elements, or main group elements. Elements can be further classified as metals, metalloids and nonmetals. Metals: solids at room temperature (except mercury) shiny good conductors of heat and electricity malleable (can be shaped) ductile (can be drawn into wires) Most metals have a high melting point and density. Common metals include aluminum, gold, platinum, silver, tin and iron Metals readily combine with nonmetals to form compounds but rarely combine with other metals. Some metals are found in their free states, like copper, gold and silver. Some metals are mixed to form homogeneous mixtures called alloys, like brass, bronze, steel and coinage metals. Nonmetals: • are not shiny • have fairly low melting points and densities • are poor conductors of heat and electricity Nonmetals combine with each other to form molecular compounds and with metals to form ionic compounds. Common nonmetals found naturally include carbon, nitrogen, oxygen and sulfur. Nonmetals Metalloids have properties between metals and nonmetals. These elements are positioned diagonally on the Periodic Table separating the metals and nonmetals. Certain metalloids, like boron, silicon and germanium are the materials used in electronic devices. Elements tend to be reactive and combine with other elements to form compounds. It is rare to find elements in nature in their pure forms. Gold, silver and platinum (the noble metals) have low reactivity and are found uncombined in nature. Seven elements exist as diatomic molecules. These molecules contain exactly two atoms A compound is a substance containing two or more elements that are chemically combined in a definite proportion by mass. Compounds, unlike elements, can be decomposed chemically into simpler substances. Compounds fall into two general types, molecular and ionic. Molecular compounds are held together by covalent bonds and ionic compounds are held together by attractive forces between the positive and negative charges. Molecules are the smallest unchanged individual unit of a compound formed by two or more atoms. A molecule cannot be further divided without destroying its identity. Ions are charged atoms or groups of atoms. Ions can be positively or negatively charged. A positively charged ion is called a cation. A negatively charged ion is called an anion. Chemical formulas are abbreviations for compounds. A chemical formula contains the symbols and ratio of the atoms of the elements in a compound. The numbers after the element symbols are subscripts. 1. The formula of a compound contains the symbols of all the elements in the compound. 2. When a formula contains only one atom of an element, the number 1 is not needed as a subscript. 3. When a formula contains more than one atom of an element, the number is indicated by a subscript written after the symbol. Writing Formulas of Compounds 4. When the formula contains more than one of a group of atoms that occur as a unit, parentheses are placed around the group and the number of units is represented by a subscript outside the parentheses. 5. Formulas show the number and kind of atoms in a compound, but not the connectivity of the elements. The Law of Definite Composition states that a compound always contains two or more elements chemically combined in a definite proportion by mass. The Law of Multiple Proportions states that atoms of two or more elements may combine in different ratios to produce more than one compound. LECTURE 5 Properties are classified as either physical or chemical. Physical properties are inherent characteristics that can be determined without altering the composition. Examples include: color taste odor physical state density melting point boiling point Chemical properties describe the ability of a substance to either undergo a reaction with another substance or to decompose. Physical changes are changes in physical properties (such as density) or changes in states of matter without a change in composition. No new substances are formed during a physical change! Sawing wood is a physical change – the wood changes shape, but the resulting pieces are still wood! In a chemical change, new substances are formed that have different properties and composition from the original material. When copper metal (Cu) is heated in air, the shiny metal turns black as copper(II) oxide is formed on the surface. A physical change often accompanies a chemical change. To succeed in chemistry, you must learn to solve complicated problems. Read the problem carefully and determine what is known and desired. Use units! READ Determine the unit relationships needed to solve the problem. Set up the problem so that the units cancel correctly. PLAN Do the math. Make sure the answer contains the proper units and significant figures. CALCULATE Check the answer — is it reasonable Energy is the capacity of matter to do work. There are many types of energy including mechanical, chemical, electrical and nuclear energy. Potential energy (PE) is stored energy, the energy an object possesses due to its position. A ball located 20 feet above the ground has more PE than when it is located 10 feet above the ground. A diver poised on a diving board has a large amount of PE. When the diver leaves the board, the energy is converted. All chemical reactions either absorb or release energy. Chemical changes can produce different kinds of energy, like electrical energy in a lead storage battery or heat and light when fuel undergoes combustion. Chemical changes can also use energy, such as the electricity used to decompose water or the solar energy used by plants during photosynthesis. Energy can be changed from one form to another or from one substance to another. The Law of Conservation of Energy states that energy can be neither created nor destroyed. When water decomposes, energy is absorbed by the system so H2 and O2 have higher potential energy. When hydrogen (H2) is used as a fuel, energy is released and water (the product) has lower potential energy. Energy comes from many sources, including petroleum, coal and woody plants, all derived from the sun. We use petroleum deposits in the forms of gasoline and natural gas. Petroleum is composed of hydrocarbons, compounds containing only carbon and hydrogen in differing ratios. Natural gas consists mainly of methane (CH4) with small amounts of ethane, propane and butane mixed in. LECTURE 6 Properties of Electric Charge 1. Charge may be either positive or negative. 2. Opposite charges (positive and negative) attract while like charges (i.e. negative and negative) repel. 3. Charge may be transferred from one object to another, by contact or induction. 4. The force of attraction between charges (F) is related to the distance between charges by: Michael Faraday: English scientist who discovered electrolytes (compounds that conduct electricity when dissolved in water). Faraday also discovered that some compounds decompose in water into their elements. These charged elements are called ions. A light bulb glows when ions are present in a saltwater solution when current is passed through it. These elements were attracted to either negatively or positively charged electrodes in the solution, meaning they were no longer neutral. Atoms are built from three subatomic particles... Protons Electrons Neutrons How was the electron discovered? Crookes Tube Mysterious Cathode Rays Sir William Crookes 1832-1919 Cathode rays were invisible solid particles coming from the metal cathode J. J. Thomson Experiment Discovered the electron (negatively charged nearly massless particle. J. J. Thompson showed that Cathode Rays were negatively charged and nearly massless electrons. The charge to mass ratio of a cathode ray is independent of the composition of the cathode or any residual gas in the tube. How was the proton discovered? Anode Rays The charge to mass ratio of anode rays is dependent on the composition of the cathode and that of any residual gas. Anode rays are positively charged When hydrogen gas was introduced into a Crooke tube an anode ray was produced with the highest positive charge to mass ratio. This smallest positively charged particle was the proton (H+). Electrons A particle with negative electrical charge Protons A particle with positive electrical charge (assigned a relative charge of +1). Protons have a much larger mass How was the nucleus of the atom discovered? Thomson’s plum pudding model: An atom is a pudding of positive charge with negatively charged plums J. J. Thomson knew the atom contained electrons and positively charged protons How was the Structure of the Atom Discovered? 1871-1937 Lord Rutherford The Rutherford Scattering Experiment Shoot alpha particles (2 4 He) from radioactive Polonium at a thin gold foil. In 1911, Ernest Rutherford established the nuclear model of the atom by bombarding gold atoms with α particles. How was the neutron discovered? James Chadwick (1891 – 1974) English scientist Rutherford to Chadwick: Look for a neutral proton Chadwick discovered the neutron Chadwick built a neutron cannon neutron had no charge and a mass almost the same as the proton. Protons and neutrons are located in the nucleus. Electrons are dispersed throughout the remainder of the atom (mainly open space). Neutral atoms contain the same number of protons and neutrons to maintain charge balance. Atoms are composed of three smaller, subatomic particles: electrons, protons and neutrons. Atomic Number: Number of protons in the nucleus of an atom. Atomic numbers for every element are above the element’s symbol in the periodic table. The atomic number determines the identity of the atom. LECTURE 7 The Nuclear Model of the Atom Atoms consist of mostly empty space except for a relatively tiny nucleus 2. All of the positively-charged and neutral subatomic particles, and most of the mass of the atom, are in the nucleus 3. The mostly-empty-space portion of the atom is more than 10,000 times the diameter of the nucleus 4. Tiny negatively-charged electrons exist in the mostly empty space and balance the positive charge on the nucleus. The unit of mass is given in atomic mass units (amu). One amu is defined as 1/12 of the mass of an atom of carbon with 6 protons and 6 neutrons in its nucleus. 1 amu = 1.6605 x 10-24 g Atom type determined by number of protons in nucleus # of protons = atomic number = Z Atomic numbers are integers that increase through the periodic table Atomic Number Every atom of any particular element has the same number of protons. The number of protons (+ charges) in an atom of an element is the atomic number, Z. For electrically neutral atoms the number of electrons (- charges) are equal to the number of protons (+ charges). Isotopes All atoms of any particular element are not identical; some have more or less mass than others. Since the number of protons and electrons in a neutral atom of an element are fixed, there must be variation in the number of neutrons that causes this difference in mass. Atoms of the same element that have different masses— different numbers of neutrons—are called isotopes. Mass number = number of protons + neutrons # protons = Atomic Number = 24 Element: Cr Atomic Mass Percentage Abundance of Some Natural Isotopes Since most elements are a mixture of isotopes, the atomic mass for an element is the weighted average of all naturally occurring isotopes of the element. Example: The atomic mass of Cu is 63.546 amu. Cu exists as 2 major isotopes, Cu-63 and Cu-65. Cu-63 is more abundant, as the atomic mass is very close to 63 amu. Calculating average atomic mass: Sum of the atomic mass of each isotope multiplied by its % abundance. Measuring Cu isotope abundances by using mass spectrometry. (62.9298) x (0.6909) + (64.9278) x (0.3091) = 63.55 amu LECTURE 8 Chemical nomenclature is the systematic naming of chemical compounds. Common names are historical names of compounds which are not based on systematic rules. Example N2O Common name: nitrous oxide (laughing gas) Systematic name: dinitrogen monoxide Common names reveal nothing about the physical and chemical properties of a compound. Common names are often used because systematic names are too long and technical for everyday use. Chemists prefer systematic names that precisely identify the chemical composition of compounds. A charged particle, called an ion, can be produced by adding or removing electrons from a neutral atom. A neutral potassium atom contains 19 p+ and 19 e-. Any neutral atom that loses an electron forms a cation. Cations are named the same as their parent atoms. Any neutral atom that gains an electron forms an anion. A neutral chlorine atom contains 17 p+ and 17 e–. A chlorine ion can be formed by adding 1 e–. This gives an ion with 17 p+ and 18 e-, Cl–. Cl + e– Cl– Change the element ending to ide Atom Name Anion Name F Fluorine F- Fluoride ion Cl Chlorine Cl- Chloride ion Br Bromine Br- Bromide ion I Iodine I- Iodide ion O Oxygen O2- Oxide ion N Nitrogen N3- Nitride ion Metals form cations Charge = Group Number Group 1A and H: +1 Group 2A : +2 Group 3A (Al): +3 Nonmetals form anions Group 7A: -1 Group 6A: -2 Group 5A: -3 Charge = 8 - Group Number Ionic compounds: contain both a cation and an anion. Example NaCl (table salt) Ionic compounds must have a net charge of 0. The sum of the charges of the cations and anions in an ionic compound equal 0. Compound Ions Least Common Multiple Sum of Charges Compound Formula Sodium bromide 6.3 Writing Formulas from Names of Ionic Compounds Rules for Writing Formulas for Ionic Compounds 1. Write the metal ion formula followed by the nonmetal ion formula. 2. Combine the smallest whole numbers of each ion to provide an overall charge equal to zero. 3. Write the compound formula for the metal and nonmetal, using subscripts determined from Step 2 for each ion. Write the chemical formula for magnesium oxide. Step 1: Write the formula for the metal and nonmetal ions. Step 2: Determine the number of each ion that will provide a net charge of 0. Step 3: Write the correct formula. Binary compounds can be either ionic or molecular. Binary compounds containing a metal which forms only one cation. Ionic binary compounds can be further subdivided: Binary compounds containing a metal which can form multiple cations. Binary compounds containing a metal which forms only one cation. By convention, the cation is written/named first followed by the anion. Transition metals can often form more than one type of cation. Example Cu can exist as either Cu+ or Cu2+ To specify the cation charge in a compound, a Roman numeral is placed directly after the metal in the compound name. Example iron(III) chloride FeCl3 Fe3+ When a metal forms only one cation (ie Na+), there is no need to use Roman numerals. Rules for Naming Compounds Involving Metals that Could Form Multiple Charges 1. Write the cation name. 2. Write the cation charge in Roman numerals in parentheses. 3. Write the root of the anion and use the – ide suffix. Exception: for metals that only form two cations, a Latin root with either an – ous or – ic suffix can also be used. Name the compound CrCl3. Step 1: Name the cation. Recognize Cr is a transition metal and can have more than one possible charge. Step 2: Use the charge of the nonmetal (Cl–) to help choose the Roman numeral and name for Cr. To charge balance, Cr must be chromium(III) Step 3: Name the anion chloride Molecular compounds contain two nonmetals. Rules for Naming Molecular Compounds o 1. Write the name for the first element, including the appropriate prefix ( mono is rarely used). o 2. Write the name for the second element, including the appropriate prefix and - ide ending ( mono is used for the 2nd element). Prefix Number Prefix Number mono 1 hexa 6 di 2 hepta 7 tri 3 octa 8 tetra 4 nona 9 penta 5 deca 10
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