Chemistry Unit #1 Matter Chemistry the scientific study of matter, its properties, and interactions with other matter and with energy a general term for many specialized fields Analytical chemistry concerned with the composition of matter development of advanced equipment Biochemistry chemistry of living things structure and processes of organisms section 1.2 Chemical engineering application of chemistry to make useful things or solve a problem Inorganic chemistry materials with non-biological origins Organic chemistry chemistry of compounds with carbon bonded to hydrogen Physical chemistry application of physics to chemistry in terms of energy and speed section 1.2 Matter a term for any type of material anything that has mass and takes up space Classification of Matter by... 1. size of matter 2. state of matter 3. type of matter section 1.2 Classification of Matter by SIZE Cosmic - matter that requires special instruments to view at great distances Macroscopic - can be seen without magnification shape color Microscopic - can be seen with a microscope dots in newsprint cell structures Submicroscopic - can t be seen with a microscope atom - a submicroscopic particle that is the smallest unit that retains all the properties of that element section 1.2
Classification of Matter by STATE vary based on temperature and the nature of the material Bose-Einstein Condensate BEC named after Bose and Einstein who described this in the 1920's first formed in 1995 by Cornell and Wieman occurs at very low temperature...billionths of a Kelvin over absolute zero atoms lose so much energy that they clump together Solid - definite shape and volume section 3.1 particles closely arranged and very tightly packed - incompressible particles in constant motion limited to vibrations expands slightly when heated Most: crystalline - regular 3-D repeating patterns Few: amorphous - no internal pattern (plastic, glass, gels, rubber) Liquid -definite volume, no definite shape liquid particles have some freedom of motion particles can move past one another - flows ( fluidity - the ability to flow) expands slightly when heated take the shape of the container section 3.1 Viscosity - measure of the resistance of a liquid to flow determined by: 1. type of attractive forces within the material stronger forces = higher viscosity 2. shape and size of the particles larger mass & slower velocity = higher viscosity 3. temperature kinetic energy increases=lower viscosity section 13.3 attractive forces limit the range of motion and keep the particles closely packed - compression is minimal
Surface Tension energy required to increase the surface area of a liquid stronger attractive forces = higher surface tension water has a high surface tension Surfactants - compounds that lower the surface tension section 13.3 Capillary action Two forces: 1. Cohesion force of attraction between identical molecules 2. Adhesion force of attraction between different molecules the force between water and the surface section 13.3 soaps and detergents Meniscus concave shape at the top of liquid in a glass container results from the water having more adhesive force and less cohesive force read the value to the bottom of a concave meniscus section 13.3 section 3.1 Gas - no definite volume or shape flows expands to fill the volume of the container particles are very far apart - compressible vapor - used to describe the gaseous state of a substance that is a solid or liquid at room temperature water vapor mercury vapor
` Plasma similar to a gas with some differences the electrons are free from the nucleus in a plasma, but in a gas the electrons are still associated with the nucleus since the electrons are removed, the nuclei are charged...ions natural plasmas: stars, Northern lights and ball lightning man-made plasmas: fluorescent lights, neon lights Phase Diagrams a graph of the pressure vs. temperature shows which phase a substance exists under different conditions different for every substance phase changes melting, freezing, vaporization (boiling), condensation (liquefaction), sublimation, deposition section 13.4 triple point where all three phases of matter can coexist all six phase changes can occur at this set of conditions critical point indicates the critical pressure and temperature above which the substance cannot exist as a liquid section 13.4 Melting point temperature at which the solid and liquid phases can coexist Vaporization (boiling) the process of particles escaping the attractive forces of a liquid to enter the gas phase
Evaporation process when vaporization occurs only at the surface process require energy, but can happen at any temperature it happens faster at higher temperatures method your body uses to control temperature in a closed container, the vapor collects above the liquid and exerts a vapor pressure on the surface of the liquid boiling point - the temperature where the vapor pressure is equal to atmospheric pressure Sublimation process by which a solid changes directly to a gas solid CO 2 DRY ICE Deposition - opposite process Condensation the gas loses energy and its particles begin to interact Freezing point temperature at which the liquid & solid phases can coexist Classification by TYPE of matter Pure substances matter with uniform and definite composition cannot be altered by physical methods Elements substance that is composed of only 1 type of atom building blocks of matter cannot be further separated by chemical or physical methods 91 naturally occurring elements scientists developed the rest of the elements U is the largest natural element section 3.4 Chemical Symbols Shorthand notation for the name of each element First letter is always a CAPITAL LETTER Second letter (if there is one) is always lowercase Symbols are the same worldwide. Origins of the Names Greek, Latin or German names based on properties Location or Scientist of Discovery Commemoration of famous scientist section 3.4
Compounds substance that is composed of two or more elements that are combined chemically Properties of the compound are generally very different from the elements that make it. Chemical Formulas Formulas show the symbols and the ratio of the elements in the compound. Subscripts tell the number of each element in the compound ex. C 12H 22O 11 section 3.4 Elements you should know: List will be provided. Compounds you should know: water H 2O glucose C 6H 12O 6 sucrose C 12H 22O 11 carbon dioxide CO 2 methane CH 4 ammonia NH 3 section 3.4 Mixtures of Matter combination of two or more pure substances in which each retains its individual chemical properties Heterogeneous mixtures not uniform in composition individual substances remain distinct Each sample will have parts in different amounts example: soup, sand and water section 3.3
Homogeneous mixture constant, uniform composition Each sample will have the parts in the same ratio example: tea, air, saltwater, antifreeze section 3.3 Solution - another name for homogeneous mixture alloy a solid-solid solution, usually of two metals steel and brass 2. Distillation used for homogeneous liquid-liquid mixtures based on differences in boiling points Separating mixtures mixtures are a physical combination separation techniques use differences in physical properties 1. Filtration used for heterogeneous solidliquid mixtures porous filter paper traps the solid as the liquids pass through section 3.3 section 3.3 section 3.3 3. Crystallization pure solid forms from a solution
section 3.3 4. Chromatography separates components based on relative attraction to two separate phases (mobile and stationary) various components move at different rates mixtures Matter separate physically substances section 3.3-3.4 chemically heterogeneous homogeneous elements compounds combine Properties of Matter Physical property - characteristic that can be observed or measured without changing the sample s composition 2 types of physical properties Extensive properties - depends on the quantity of matter mass, length, volume Intensive properties - independent of the quantity present melting point boiling point density Very useful for substance identification section 3.1 Physical changes alter a substance without changing its composition changes in state - all transitions like melting, freezing, boiling and condensing Chemical property - ability of a substance to combine with or change into one or more substances Can only be determined by changing the substance Example: iron combines with oxygen to form iron (II) oxide (RUST) Note: each substance has a UNIQUE set of chemical and physical properties. section 3.1
Chemical changes a process that involves one or more substances changing into a NEW substance known as a chemical reaction new chemical: different composition and properties Evidence of a chemical reaction 1. color changes 2. energy changes - gets cold or hot 3. odor changes 4. precipitate - solid formed from a mixture of solutions 5. gas produced - bubbles 6. irreversible process 7. new properties section 3.2 VOLCANO.MOV Quantifying matter Units of measurement Units are important because they tell us what the number represents. section 2.1 SI Units revised metric system base units - defined unit based on an object or event in the physical world SI base units (used in this course) Time second s Length meter m Mass kilogram kg Temperature kelvin K Amount of substance mole mol Electric current ampere A Luminous intensity candela cd Metric prefixes same set of prefixes used with ALL units prefixes are abbreviated with units Prefixes Used with SI Units Prefix Symbol EX: equivalence in meters giga- G 1 Gm = 1 000 000 000 m mega- M 1 Mm = 1 000 000 m kilo- k 1 km = 1000 m deci- d 10 dm = 1 m centi- c 100 cm = 1m milli- m 1000 mm = 1 m micro- 1 000 000 m = 1 m nano- n 1 000 000 000 nm = 1m section 2.1 Derived Units units defined by a combination of other units most of the units that are used Volume space occupied by an object cubic meter m 3 L - liter 1000 ml = 1 L 1 ml = 1 cm 3 = 1 cc (medicine) Density - ratio of the mass of an object to its volume D = = = = section 2.1 density of water = 1.00 g/ml
Temperature a measure of the intensity of the heat in an object Temperature measurements Celsius - more convenient than Kelvin based on the freezing point 0 o C and boiling point 100 0 C of water Kelvin - K - SI base unit for temperature positive temperature scale K = 0 C + 273 0 C = K - 273 water freezes at 273 K and boils at 373 K o F = (1.8 x o C) + 32 o C = o F - 32 1.8 section 2.1 Convert 56 o C to K and o F Problem Solving Process THE PROBLEM read the problem identify the unknown:what will you be solving for? ANALYZE THE PROBLEM identify and list the given values gather information that will help you plan the steps that you will follow using an equation isolate the unknown substitute into the equation SOLVE FOR THE UNKNOWN without an equation start with the known use dimensional analysis EVALUATE YOUR ANSWER is the answer reasonable? check for units and significant figures!
Dimensional Analysis problem-solving method use the units to solve the problem in a series of steps conversion factors set the transition to each new unit conversion factor ratio of equivalent values used to express the same quantity in different units 1 gal = 4 qt ratio always equal to 1 unit changes without changing the quantity section 2.2 Conversion factors from English standard to Metric memorize one for mass, length and volume 1 lb = 453.59 g 1 kg = 2.2 lb 1 in = 2.54 cm 1 km = 0.621 mi 1 gal = 3.78 L 1 qt = 946 ml Example: Convert 3.51 in to m section 2.2 Example: A race car travels at 225.5 mph. Convert this to m/s. Scientific notation used to express large and small numbers 3.45 x 10 6 section 2.2 Example: A physician has ordered 325 mg of atropine, intramuscularly. If atropine were available as 0.50 g/ml of solution, how many cc's would you need to give? coefficient: exponent: a number whole # between 1 & 10 equal to the # of decimal places moved Exponent original number > 10 positive exponent original number < 1.negative exponent
Example: 504,000,000,000,000,000,000 m establish a coefficient 5.04 x 10 20 m move decimal & set exponent Calculating with Scientific Notation Use a calculator with exponential notation. section 2.2 3.11 x 10 3 g / 9.02 x 10-1 L = 4.53 x 10-8 m x 3.33 x 10-2 m = 4.56 x 10-15 kg 4.56e - 15 THIS LOOKS LIKE THAT on the calculator Reliability of measurements Accuracy how close a measurement is to the accepted value depends on the instrument Precision how close a series of measurements are to one another the repeatability of the measurements depends on careful, skillful measuring techniques section 2.3 Significant figures all of the digits known precisely from the instrument plus one carefully estimated digit Used to indicate the precision of measurements. section 2.3 a) 4.5 cm b) 4.58 cm c) 3.0 cm
section 2.3 Significant Figures Counting Rules A number is a significant figure if it is: 1. a non-zero digit 24.7 m 3 sig figs 714 m 3 sig figs 2. a zero between nonzero digits (between sig figs) 7003 g 4 sig figs 1.5001 g 5 sig figs 3. a zero at the end of the measurement and to right of the decimal 7.000 L 4 sig figs 4. any digit written in scientific notation 5.70 x 10-3 g 3 sig figs A number is not significant if it is: 1. a zero at the beginning of a decimal number 0.001 s 1 sig fig 0.01035 s 4 sig figs 2. a zero at the end of the measurement and to the left of the decimal 7000 m 1 sig fig 3300 m 2 sig figs Counting numbers and defined constants have an infinite number of significant numbers. 9 players 1 in = 2.54 cm section 2.3 Rounding off numbers section 2.3 results from calculations involving measurements must be rounded to the correct number of significant figures Simple rounding rules if the digit to the immediate right of the last significant figure is less than 5, do not change the last significant figure Round 4.38291625 m to 3 sig figs. 4.38291625 last sig fig is unchanged 4.38 m if the digit to the immediate right of the last significant figure is 5 or greater, round up the last significant figure Round 4.38291625 m to 4 sig figs. 4.38291625 last sig fig is changed 4.383 m Round to three sig figs: 4039.472 m 0.0098735867 g 2901456899 L
Significant Figure Rounding Rules For: Addition and Subtraction Round to the least precise decimal place 5.2 m Precise to first decimal place +1.375 m Precise to third decimal place 6.575 m Round to least precise place section 2.3 Answer = 6.6 m 3.44 12040 5.01 + 317 + 50 For: Multiplication and Division Round to the least number of sig figs 7.55 m Measurement has 3 sig figs x0.34 m Measurement has 2 sig figs 2.567 m 2 Round answer to 2 sig figs Answer = 2.6 m 2 3.500 cm x 2.95 cm= section 2.3 0.l12 g 2.0 ml = Representing Data Graphing - helps reveal if a pattern exists graph - visual display of data Types of graphs: 1. Circle graph also known as a pie chart useful for showing parts that add up to 100 % each piece of data has its own wedge section 2.4 2. Bar graph useful for showing how a quantity varies over time, location or temperature each piece of data has its own bar section 2.4
3. Line graph most commonly used in chemistry useful for showing the relationship between and independent and dependent variable independent variable x-axis dependent variable y-axis Line cannot pass through all of the scattered points, so as many points fall above the line as below it. (called the best-fit line) section 2.4 Interpreting graphs Slope of a line graph: 1. the best-fit line is straight, the variables have a linear relationship positive slope indicates a direct proportion negative slope indicates an inverse proportion 2. the best-fit line is curved variables have a nonlinear relationship section 2.4 Reading data from a graph Interpolation - reading values between measured points Extrapolation - estimate values beyond measured points by extending the line section 2.4 Percent Error % error = accepted value - experimental value accepted value x 100