Introduction to General Chemistry

1. Introduction to General Chemistry Ch. 1.1- 1.5

2. What is Chemistry? • Chemistry is the study of properties of substances and how they react • Chemical substances are composed of matter • Matter is the physical material of the universe; anything with mass that occupies space is matter • Matter can take numerous forms • Most matter is formed by unique arrangements of elementary substances called elements

4. Elements, Compounds, and Molecules • An element can easily be defined as a substance that can not be broken down into simpler substances • Millions of different materials in the world, all comprised of some combination of only 118 elements • Similar to how the alphabet combines 26 letters to yield hundreds of thousands of words, elements bond in unique arrangements to give different molecules • Molecules agglomerate to yield compounds

5. Molecules Are Comprised of Uniquely Arranged Atoms. Different Molecules Have Different Properties. O O O H H H H

6. Small Molecular Differences Can Yield Vastly Different in Terms of Biological Interactions acetaldehyde (hangover) acetic acid carbon dioxide ethanol BLINDNESS!!! methanol

7. Small Molecular Differences Can Yield Vastly Different in Terms of Biological Interactions Relief of Morning Sickness Severe Limb Defects

8. Small Molecular Differences Can Yield Vastly Different in Terms of Biological Interactions C6H12O6 C6H12O6 Technically, glucose and fructose are the same. So is high fructose corn syrup really that bad for you???

9. Small Molecular Differences Can Yield Vastly Different in Terms of Biological Interactions

10. The Properties of Molecules Differ Vastly from those of the Atoms That Comprise Them Na (sodium metal) Cl2 (chlorine gas) Na+Cl-

11. Different atomic arrangements can change physical properties Carbon (graphite vs diamond)

12. Spatial Dimensions of Compounds Can Alter Properties Gold Nanoparticles Bulk Gold 5 nm 50 nm

13. Spatial Dimensions of Compounds Can Alter Properties 2 nm 12 nm CdSe quantum dots

14. Phases of Matter: Solids, Liquids and Gases Solids • Atoms tightly bound • Fixed volume and shape (does not conform to container) • A chemical is denoted as solid by labeling it with (s) S(s)

15. Phases of Matter: Solids, Liquids and Gases Liquids • Atoms less tightly bound than solids • Has a definite volume, but not definite shape (assumes the shape of its container) • Denoted by (L) H2O (L)

16. Phases of Matter: Solids, Liquids and Gases Gases • Free atoms • No shape, no definite volume • Can be expanded or compressed (like engine piston) • Denoted by (g) ; ex. O2 (g)

17. Qualitative and Quantitative Analysis • In chemistry, the scientific method is used to investigate scientific phenomena & acquire new knowledge • Empirical evidence is gathered which supports or refutes a hypothesis • Empirical evidence is either quantitative or qualitative • Quantitative data is numerical, and results can be measured • Qualitative data is NOT numerical, but consists of observations and descriptions

18. Quantitative and Qualitative Analysis A + B C Quantitative data • How much C is formed? • How efficient is the reaction? • What is the rate of the reaction? Qualitative data • What color is it? • Is it solid, liquid, gas? • How does it smell?

19. Units • Internationally accepted system of measurements is called the SIunit system • Quantitative measurements are represented by a: NUMBER and a UNIT • A unit is a standard against which a physical quantity is compared physical quantity • Temperature is measured in Co, Ko,orFo • Currency is measured in $USD • Distance is measured in meters, miles, ft, etc. • Time is reported in seconds, minutes, hr, etc.

20. SI Unit System: The Units of Physical Science

21. Greek Prefixes • Prefixes indicate powers of 10 • ex. k= 103; 5 kg = 5 x (103)g

22. A Review of Scientific Notation Scientific notation indicates a factor (F) multiplied by a power (n) of 10 F x 10n (1 < F < 10) • Important: All integers end with a decimal point, even though it is not commonly written (1  1. ) • If no factor is shown, assume there is a 1. in front of powers of 10: 102 = 1. x 102 10-7 = 1. x 10-7 • For every positive power of 10, shift the decimal 1 place to the right, add a zero for each place 102 = 1. x 102 = 100. 105 = 1. x 105 = 100000.

23. A Review of Scientific Notation • Convert to scientific notation • 15 • 125.3 x 102 • 0.003003 • Convert to standard notation • 3.4912 x 104 • 8.971 x 10-3 • 6.50 x 100 For all non integers, simply shift the decimal. 2.5 x 105 = 250000. 1.8773 x 108 = 187730000. For negative exponents, shift the decimal left. All values less than 1 have negative exponents. 7.141 x 10-2 = .07141 3.867 x 10-7 = .0000003867

24. Multiplying and Dividing Exponents (Review) • When multiplying powers of 10, the product is the sum of the powers • 102 x 105 = 10 2+5 = 107 • When dividing powers of 10, subtract • 102 / 105 = 10 (2-5) = 10-3 • Powers of powers require multiplication • (102)5 = 1010

25. Group Work Convert the following values to gramsin proper scientific notation. • 421.4 kg • 1170.1 mg • 481 µg

26. Why Are Units Important? Example #1 • In 1999, NASA lost the $125M Mars Orbiter System. • One group of engineers failed to communicate with another that their calculated values were in English units (feet, inches, pounds), and not SI units. • The satellite, which was intended to monitor weather patterns on Mars, descended too far into the atmosphere and melted.

27. Why Are Units Important? Example #2 • In 1983, an Air Canada Plane ran out of fuel half way through its scheduled flight. Why? • Airline workers improperly converted between liters and gallons. • Luckily, no one died.

28. Why Are Units Important? Example #3 • A case was reported in which a nurse administered 0.5 g of a sedative to a patient. • The patient died soon after • The patient should have only received 0.5 grains (≈ 0.033 g) but the units were not listed. That was the equivalent of 8 doses!!

29. Derived SI Units: VOLUME • Many measured properties have units that are combinations of the fundamental SI units • Volume: defines the quantity of space an object occupies; or the capacity of fluid a container can hold • expressed in units of (length)3 or Liters (L) • 1 L is equal to the volume of fluid that a cube which is 10 cm on each side can hold 10 cm V =(10 cm)3= 1000 cm3 10 cm 1 L = 1000 cm3 10 cm 1000 mL = 1000 cm3 mL = cm3

30. Derived SI Units: DENSITY Ex. If you had equal volumes of bricks and feathers, the masses would be very different. THE DENSITY OF WATER IS All matter has mass, and must therefore occupy space. Densitycorrelates the mass of a substance to the volume of space it occupies. Density = mass per unit volume (mass/volume). Different materials have different densities. You must use density to convert between mass and volume

31. Group Work • A cubic container that is 25 cm on each side is filled with ethanol. The density of ethanol is 0.79 g/mL. • What is the volume of ethanol in the cube in mL? L? • What is the mass of ethanol in g? Give answers in scientific notation!!

32. Derived SI Units: ENERGY • What is Energy? • Energy is defined as the capacity to perform “work” • How do we define work? • Work is defined as the action of applying a force acting over some distance. Work can not be done if no energy is available. • In SI units, we use the unit Joule(J) to represent energy.

33. Conservation of Energy Energy is never created or destroyed, merely converted between forms and transferred from place to place. The total energy of the universe is finite.

34. Conservation of Energy • Ex. Power plant

35. Forms of Energy • Energy comes in many forms and can be converted from one form to another. Some examples are given: • Chemical Energy • Energy stored in chemical bonds (e.g. gasoline, coal, etc.) that can be released by chemical reaction, typically combustion (fire) • Heat Energy (thermal energy) • Heat is defined as energy flow between bodies of matter resulting from collisions of molecules or random motions of electrons.

36. Forms of Energy • Mass Energy • Energy and mass are interchangeable. During a fusion reaction (e.g. stars), mass is lost and energy is formed. This mass appears as energy according to the following: where m is the change in mass (in kg), c is the speed of light, and E is the energy released (J). This is the basis of nuclear power. • Kinetic Energy • Energy of motion (e.g. a moving car). An object with mass m, moving at a velocity V (meters/sec) has kinetic energy:

37. Forms of Energy • Potential Energy • Potential energy corresponds to energy that is stored as a result of the position of mass in a field. • If a mass m is held at a height h (meters) above the ground, assuming a gravitational acceleration of 9.8 m/s2(g), its potential energy is: • If the object is dropped, it loses potential energy. However, it speeds up as it falls, so its kinetic energy increases equally (conversion).

38. Forms of Energy • Electrical Energy • Energy resulting from electric current, the movement of electrons through a conductive circuit. Electrical energy is a type of potential energy. For a charge q (coulombs, C) moving across a voltage, V • Light/Radiation • The energy of a wave of light is calculated as the product of Planck’s constant, h (J s), and the wave frequency, ν (1/s)

39. Power • It is often necessary to express the rateof energy usage. This is called power. • Typically, we speak in terms of energy per second. In SI units, a joule per second (J/s) is known as a watt (W).

40. Temperature • Temperature: a measure of the tendency of a substance to lose or absorb heat. The value of temperature is dictated by the speed of the molecules in a system. • Faster molecules = Higher Temperature • Temperature and heat are not the same. • Heat always flows spontaneously from bodies of higher temperature to those of lower temperature

41. Temperature • When performing calculations in chemistry, temperature must always be converted to Kelvin (K) units(unless otherwise stated). • The lowest possible temperature that can ever be reached is 0K, or absolute zero. At this temperature, all molecular motion stops. • To convert temperatures to the Kelvin scale: oK : oC + 273.15