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Chapter 1

Chapter 1. AP Chemistry Milam. 1-1 matter and energy. Matter takes up space and has mass, so anything besides light, energy, forces Mass is how much matter something consists of Energy is the ability to do work (useless definition) Kinetic energy ½ mass * velocity 2. 1-1.

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Chapter 1

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  1. Chapter 1 AP Chemistry Milam

  2. 1-1 matter and energy • Matter takes up space and has mass, so anything besides light, energy, forces • Mass is how much matter something consists of • Energy is the ability to do work (useless definition) • Kinetic energy ½ mass * velocity2

  3. 1-1 • Potential energy is stored energy and there are multiple kinds of potential energy (nuclear, gravitational, chemical…) • Exothermic is a process that releases heat (surroundings get warmer) • Endothermic is a process that absorbs heat (surroundings get cooler)

  4. 1-1 • Laws of conservation – In most situations, both energy and mass are conserved, meaning if you start with a certain amount of energy or mass you end up with that much • However, mostly in nuclear reactions, energy and mass may interchange according to the equation Energy = mass * (speed of light)2 • E= mc2

  5. 1-2 Chemistry – A Molecular View of Matter • Chemistry studies very small objects (atoms and molecules), but can do so on a large scale and then relate this back to a small scale • Democritus was one of the first people to postulate the idea of atoms, although he performed no experiments • John Dalton did perform experiments….

  6. 1-2 • Dalton discovered that: • Every element is made up of atoms • All atoms of an element have identical properties • Atoms are conserved • Compounds form in small whole-number ratios • The relative # and type of atoms are constant in a given compound

  7. 1-2 about the atom • Atoms are made of electrons, protons and neutrons • Neutrons are slightly bigger than protons and both are much larger than electrons • Neutrons and protons make up the nucleus and electrons are in the electron cloud which takes up the majority of the volume of an atom

  8. 1-2 Atomic # • Atomic # (Z) is the number of protons in an atom. • Because the atoms of an element all have the same # of protons they also have the same atomic # • You can locate an element on the periodic table using atomic # (ex. Hydrogen Z=1)

  9. 1-2 Molecules • A molecule typically consists of more than 1 atom bonded together that is stable at normal conditions and electrically neutral. • There are seven diatomic molecules, these exist in nature rather than the elemental forms (H2, N2, O2, F2, Cl2, Br2, I2)

  10. 1-3 states of matter • Solids, liquids and gases will all be developed in more detail in the future, for now be able to recognize each by the pictures at the molecular level

  11. 1-4 Chemical and Physical Properties • Chemical properties are how a substance will and will not react with other chemicals • A chemical property of sodium is that it reacts violently with water • Physical Properties are any other property • Color, density, hardness, melting point, shape, solubility….. • Ice melts at 0° C is a physical property of water

  12. 1-4 • Extensive properties depend on the amount of the material • Volume, mass, weight…. • Intensive properties are independent of amount • Color, melting point, density, chemical properties….

  13. 1-5 chemical and physical changes • Chemical changes involve a chemical reaction • Signs of a chemical change include change in color, gases or light being produced, energy being absorbed and released, forming a precipitate • Physical changes involve any change where the chemicals remain the same

  14. 1-5 • Examples of physical changes include: phase changes (melt, sublimate, condense), dissolving, breaking, ripping…. • Physical changes would not protect you from poison, chemical changes might or might not depending on the new product

  15. 1-6 Mixtures, substances, compounds and elements • Most everyday items are mixtures, there are two kinds of mixtures • Heterogeneous – where the mixture is not uniform • Ex. Salad, Legos, sand • Homogeneous (aka solution) – where the mixture is uniform • Ex. Kool-aid, alloys, air

  16. 1-6 • An element is a type of atom with a specific number of protons, an element cannot be decomposed into a simpler substance • Ex. Br2, Cu, He, Silver metal • A compound is a substance with two or more types of atoms in a fixed ratio of amounts • Ex. Water H2O, C6H12O6, H2SO4

  17. 1-6 • The law of definite proportions (law of constant composition) says that if you take a compound and compare the percent mass of each element, they will always be the same for each compound • So if you analyze pure water to be 88.9% oxygen and 11.1% hydrogen, it will always have those percentages.

  18. 1-6 • The table with elements and symbols on page 18 is worth becoming familiar with • Spend 5 minutes 2-3 times looking over the list and memorizing the elements and symbols that you aren’t already familiar with

  19. 1-7 Measurements in Chemistry • Being familiar with some of the prefixes in table 1-6 on page 19 is important, specifically • Kilo, Centi, milli, micro and nano are commonly used • Pico, mega, deci are used in some odd situations, but when those arise you can just look them up

  20. 1-8 Units of Measurement • You are probably already familiar with most of the units, but maybe you are not used to using some of the SI units • You need to know: length – meters; Mass – kilograms; Volume – liters; time – seconds; temperature – kelvins; amount of substance - mole

  21. 1-9 Use of Numbers • Scientific notation is a way of writing really small and really large numbers • The basis for it is that 108 = 100,000,000 so if I am doing a calculation and get an answer of 200,000,000 I can write 2x108 instead • Be able to convert between numerical and scientific notation

  22. 1-9 sci. not. • Ex. 140,678 into scientific notation • 1.40678 x 105 • Move the decimal until the # is between 1 and 9.9999…… • Count the number of decimal places moved and make that the exponent • Ex. 2, turn 4.57 x 10-7 into numerical notation

  23. 1-9 sci. not. • Ex. 2, turn 4.57 x 10-7 into numerical notation • 0.000 000457 • Negative exponents give numbers less than 1 and positive exponents give numbers greater than 1 • 100 = 1

  24. 1-9 sig figs • Significant figures tell you how accurate of a measurement was made • To take a reading, you read the smallest increment on the measuring tool and estimate one additional place • Once you take this reading you need to keep the amount of significant figures consistent when you make calculations from your data

  25. 1-9 sig figs • In this class we assume that all calculations are from lab data, hence, all calculations must keep significant figures intact • Significant figures are worth points on the free response sections of the test and could show up on the multiple choice as well • All digits 1-9 are always significant, 0’s depend on the situation

  26. 1-9 sig figs • 0’s that are in between two non zero digits are always significant (ex. 303 is 3 sig figs) • If there is no decimal: • All 0’s are not significant • If there is a decimal: • All 0’s to the right of the last nonzero digit are significant

  27. 1-9 sig figs • Ex. 1 - 05000 has 1 significant digit • Ex. 2 – 6470.0 has 5 sig figs • Ex. 3 – 0.004 has 1 sig fig • Ex. 4 – 0.00320 has 3 sig figs • Ex. 5 – 0.0040300 has 5 sig figs • Sometimes it helps to read from right to left with a decimal, and the last nonzero number is the last sig fig

  28. 1-9 sig figs • In addition and subtraction the significant figures are determined by the least significant place • Adding 4 + 0.78 gives 4.78, but since 4 only has a significant figure in the ones place, the answer needs to be good only to the ones place as well • 4 + 0.78 = 5

  29. 1-9 sig figs • Multiplying dividing is a lot more common, especially since conversions involve these • To maintain the number of significant figures you count the number of significant figures for each thing being multiplied or divided and the one with the least number is how many your answer has • Ex. 4.0 x 2.00 x 1.0000 = _________

  30. 1-9 sig figs • Ex. 4.0 x 2.00 x 1.0000 = _________ • 4.0 has 2, 2.00 has 3 and 1.0000 has 5 sig figs, so your answer should have 2 • 4.0 x 2.00 x 1.0000 = 8.0000 = 8.0 • Exceptions, exact numbers have infinite sig figs • Ex. Counting numbers (2 people), conversion factors (3 feet = 1 yard)

  31. 1-9 sig figs • Rounding sig figs to X places • To round 1.2345678 to 4 places, you go to the 4th place (the 4), then if the number directly to the left is 5 or larger you round up, if it is 4 or lower you round down • 1.235 • Significant figures and significant digits are the same thing and will be used interchangeably as will sig figs and sig digs

  32. 1-9 sig figs • The most important things about sig figs at this point is understanding why we use them, and being able to correctly count 0’s as significant or not. The rest will get lots of practice.

  33. 1-10 The unit factor method • This chapter is devoted to conversions, there are two ways to understand conversions • I think of conversions as multiplying by one, you take some value, and multiply it and divide by two equivalent values that are different units • You can also understand a conversion as a process and use matching units and routine to do them

  34. 1-10 • Ex. Convert 24 inches into feet • 24 inches * (1 foot) = 2 feet (12 inches) • Because 1 foot = 12 inches, you have changed the units without changing the value of it

  35. 1-10 • Ex. Convert 2 miles into feet • Step 1 write down given • 2 miles • Step 2 make the crossbow • 2 miles * ( ) ( )

  36. 1-10 • Step 3 fill in units of given on bottom and desired units on top • 2 miles * ( feet ) ( miles ) • Step 4 Fill in equivalent amounts • 2 miles * ( 5280 feet ) = 10,560 feet ( 1 mile )

  37. 1-11 Percentage • % by mass is used in chemistry and it is the same as the rest of the percentage calculations you’ve done previously, it’s just more intimidating because it involves chemicals • Always remember that percents are a part of a whole, so you will always have a part, and the whole.

  38. 1-12 Density and specific gravity • Density is mass/volume, it is an intensive property so it does not depend on amount • Density makes a good conversion factor between volume and mass • Specific density is the density of an object relative to the density of water • To get specific density you just remove the units (or divide by the density of water)

  39. 1-13 Heat and temperature • Converting between fahrenheit and celsius is almost never needed and can be looked up when it is, however, Celsius to Kelvin conversions occur all of the time • Since the lowest temperature possible is 0 K, and Celsius can be negative, to get from Kelvins to Celsius you subtract 273 and add 273 to get from Celsius to Kelvin

  40. 1-13 • You should also know the freezing and boiling points of water in all three scales • Fahrenheit 32 and 212 • Celsius 0 and 100 • Kelvin 273 and 373 • Kelvins do not include the degree symbol

  41. 1-14 heat transfer and the measurement of heat • Joules are the most common units of energy • Calories are also used, a calorie is 4.184 J, and a Calorie or kilocalorie is 4184 J. The Calorie with a big C is the kind you see on food packaging • Specific heat capacity is the amount of heat it takes to heat up 1 gram of a substance

  42. 1-14 Specific heat capacity • Q = mCt where Q is heat (J), m is mass (g), C is a specific heat constant (J/(g °C), and t is the change in temperature (°C) • C depends on the material, for water C is 4.184 J/(g °C) and for metals the specific heat is much lower because they heat up to higher temperatures with less heat

  43. 1-14 • Heat capacity is like specific heat capacity only it is for an entire object, not 1 gram of a substance. Heat capacity is often used for calorimeters and bomb calorimeters and when the mass is unknown • A calorimeter is a device that usually consists mostly of water that the temperature change can be related to find the amount of heat absorbed or released

  44. 1-14 • Sample calorimetry problem

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