1. Chemistry 400 Ch 1

2. Matter  Chemistry is the study of matter and its properties, including how it reacts.  What’s matter? Anything with mass and volume (or occupies space).  What are the 3 states of matter (or phases of matter) that chemists study? Gas, Liquid, and Solid

3. Atoms, Elements, Compounds  Atoms are the smallest particle which retains its chemical identity.  Elements are substances which can’t be broken down into simpler substances by chemical means.  Where do you find a chart of elements?

4. Elements and the Periodic Table

5. Atoms, Elements, Compounds  Elements are found in nature as elemental forms.  Most elements are found as atomic elements, but some elemental forms are molecules.  You must know the 7 diatomic elements (diatomic molecules).

6. Atoms, Elements, Compounds  Compounds are substances which are composed of more than 1 element.  Compounds can be broken down into simple substances (elements) by chemical means.  How can we represent a compound like water or table salt? Chemical Formulas

7. Pure Substances & Mixtures  Pure substances are substances which contain only 1 kind of substance.  A pure substance can either be an element like gold or oxygen gas, or it can be a pure chemical compound like NaCl or H 2 O.

8. Pure Substances & Mixtures  Pure substances can’t be separated into different substances by physical means like filtration or distillation.  Compounds can be separated or broken down into different substances by chemical means or chemical reactions.  Elements can’t be separated or broken down into different substances by chemical reactions.

9. Pure Substances & Mixtures  Mixtures contain 2 or more substances, which are physically mixed together.  There are 2 kinds of mixtures: homogeneous (also called solutions) and heterogeneous.  Homogeneous: mixture where all particles are uniformly distributed; these are solutions like air, bronze, or saltwater.  Heterogeneous: mixture where the particles are unevenly distributed; this can often be seen as different layers like oil and water.

10. Pure Substances & Mixtures

11. Properties  What are properties? They are characteristics of a substance like taste, odor, color, etc.  Properties may be classified as chemical or physical properties or they may be classified as extensive or intensive properties.  A physical property is a property which does not change the chemical’s identity.  What are some examples of physical properties?___________________________

12. Properties  A chemical property is one which does change the chemical’s identity.  What are some examples of chemical  properties?___________________________  Intensive properties are properties, like boiling point, which do not depend on the amount of the substance.  Extensive properties like volume or mass do depend on the amount.

13. Law of Conservation of Mass  Matter is conserved: it is neither created nor destroyed in chemical reactions.  What does this mean?

14. Energy  Energy is the ability or capacity of something to do (perform) work.  What types of energy are there?  As there is a Law of Conservation of Mass, there is a Law of Conservation of Energy.

15. Energy  If chemicals contain potential energy, or a stored energy, do you think that chemicals prefer to have a high potential energy or a low potential energy?  How do you think this relates to the stability of a chemical?  Is a chemical with a high potential energy more or less stable than a chemical with a very low potential energy?

16. Measurements  In the lab, we make measurements, that is we collect data with numbers and units.  The basic types of measurements with simple units are mass, length, time, amount, current, temperature, and luminosity.  Although there are many different ways of expressing these measurements, there is an approved system of units, the SI system. That way, any scientist can communicate with and understand the measurements that any other scientist makes.

17. Measurements MeasurementSI Unit & AbbrevUS Unit & Abbrev MassKilogram, kgPound, lb LengthMeter, mFeet, ft TempKelvin, KFahrenheit, o F AmountMole, molsame CurrentAmpere, amp, Asame LuminosityCandela, cdsame TimeSecond, sec, ssame

18. Metric System  If we are using very small or very large numbers, we use the metric system of prefixes.  You need to memorize and use the following: GigaMega KiloDeci CentiMilli MicroNano Pico

19. Metric System PrefixAbbrevMeaningExample GigaG1x10 9 (1,000,000,000)1GW = 1x10 9 W MegaM1x10 6 (1,000,000)1Ms = 10 6 s Kilok1x10 3 (1,000)1km = 1000m Decid1x10 -1 (0.1)1dL = 0.1L Centic1x10 -2 (0.01)1cm = 0.01m Millim1x10 -3 (0.001)1mg = 10 -3 g Micro  1x10 -6 (0.000001) 1  A = 10 -6 A Nanon1x10 -9 (0.000000001)1ns = 10 -9 s Picop1x10 -12 1pg = 10 -12 g

20. Scientific Notation  We can also express very large or small numbers in scientific notation.  Scientific notation is a type of exponential notation.  There is always just 1 nonzero digit to the left of the decimal.  Which of these is Scientific Notation? 1.4x10 -6 34.0x10 3 5.67.245379x10 8

21. Scientific Notation  You have to be able to convert normal or “floating” numbers to scientific notation and vice versa. (You will practice in the lab.)  Ex. Convert 17854 and 0.0000457 to scientific notation.  Ex. Convert 4.67x10 -6 to “floating” notation.  Many of your calculators can do this for you easily: ask me or your lab instructor!

22. Mass vs. Weight  Although you will hear many people (including chemists) use the terms mass and weight interchangeably, they are NOT the same!  Mass is independent of gravity. Your mass here is the same as your mass on Mt. Everest is the same as your mass on the moon.  Weight depends on gravity. Your weight is less on Mt. Everest and is much less on the moon!  We measure mass on lab balances. (but often just say weight)

23. Temperature Scales  Scientists all over the world use 2 temperature scales: Celsius (°C) and Kelvin (K)  The celsius scale is a relative scale.  It’s numbers were set “relative” to the boiling point (set at 100°C) and freezing point (set at 0°C) of water.  The Kelvin scale is an absolute scale.  The lowest temperature possible in our Universe (based on physics) is Absolute Zero, or 0K.  This is the temperature at which all molecular motion ceases.

24. Temperature Scales  Note that the Kelvin scale stops at 0. It can never be a negative number!  The Celsius scale can be negative. It stops at -273.15°C.  How do you convert between K and °C?  They differ by 273.15! K = °C + 273.15 OR °C = K - 273.15  Remember that the Kelvin temp is higher than the Celsius temp.

25. Temperature Scales  In the US, we also use the Fahrenheit scale, °F.  The Fahrenheit scale is also a relative scale.  It was based on our blood temperature as well as the boiling point of water.  The Fahrenheit scale may also be negative.  How do you convert between °F and °C? °F = (9/5)°C + 32 OR °C = (5/9)(°F - 32)

26. Derived Units (Compound Units)  Scientists also make measurements with derived or compound units.  These are units which contain more than 1 of the basic SI units (mass, length, time, etc.)  Volume is a derived unit as it is a compound unit of m 3.  Area is m 2.  Can you think of any?

27. Derived Units (Compound Units) Name, AbbrevSI UnitOther Common Units Area, Am2m2 Volume, Vm3m3 cm 3, mL, L, dL Density, dkg/m 3 g/cm 3, g/mL, g/L Speed, vm/skm/hr Pressure, PPa (kg/ms 2 )atm, torr, kPa Energy, E, H, GJ (kgm 2 /s 2 )kJ, cal, kcal

28. Density  Density is a common compound unit and is an intensive physical property of a substance.  It is a measure of how close the particles of the substance are. Or how tightly the substance is packed.  It is mass/volume OR d = m/V.  We commonly use g/mL or g/cm 3 (1mL = 1cm 3 ) for solids or liquids.  We usually use g/L for gases.

29. Density  Since we know that d = m/V, we can solve for d, m, or V.  If the density of a soil sample is 2.431g/cm 3, what is the mass of a 534cm 3 sample?  Density depends on the temperature, as you know if you think about water.  Usually, the density of a substance increases as it goes from the gas to liquid to solid phase (i.e. the temp decreases).  Is this true for water? What does ice do?