Introduction to General Chemistry

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

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

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

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

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

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

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

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

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

Phases of Matter: Solids, Liquids and Gases 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)

Phases of Matter: Solids, Liquids and Gases 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)

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

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

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?

Units 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,or Fo Currency is measured in $USD Distance is measured in meters, miles, ft, etc. Time is reported in seconds, minutes, hr, etc. Internationally accepted system of measurements is called the SI unit system

SI Unit System: The Units of Physical Science

Greek Prefixes Prefixes indicate powers of 10 ex. k= 103; 5 kg = 5 x (103)g GP1

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.

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.

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 (1 grain = 0.065g) but the units were not listed. That was the equivalent of 8 doses!!

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 1 L = 1000 cm3 1000 mL = 1000 cm3 mL = cm3

THE DENSITY OF WATER IS 𝟏 𝒈 𝒄𝒎 𝟑 𝒐𝒓 𝟏 𝒈 𝒎𝑳 Derived SI Units: DENSITY All matter has mass, and must therefore occupy space. Density correlates the mass of a substance to the volume of space it occupies. Density = mass per unit volume (mass/volume). Different materials have different densities. A container filled with bricks does not have the same mass as an equivalent volume of feathers! GP2 THE DENSITY OF WATER IS 𝟏 𝒈 𝒄𝒎 𝟑 𝒐𝒓 𝟏 𝒈 𝒎𝑳

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 Ex. Pushing an object along a rough surface In SI units, we use the unit Joule (J) to represent energy. 𝐽= 𝑘𝑔 𝑚 2 𝑠 2

Temperature Temperature: a measure of the tendency of a substance to lose or absorb heat. Temperature and heat are not the same. Heat always flows from bodies of higher temperature to those of lower temperature An active stove top is ‘hot’ because the surface is at a much higher temperature than your hand, so heat flows rapidly from the stove to your hand Ice feels ‘cold’ because it is at a lower temperature than your body, so heat flows from your body to the ice, causing it to melt Explain wind chill?

Temperature When performing calculations in chemistry, temperature must always be converted to Kelvin (oK) units (unless otherwise stated). The lowest possible temperature that can ever be reached is 0oK, or absolute zero. At this temperature, all molecular motion stops. To convert temperatures to celcius and Kelvin: oK : oC + 273.15 oC: 5/9 (oF – 32) GP3

Accuracy and Precision Accuracy defines how close to the correct answer you are. Precision defines how repeatable your result is. Ideally, data should be both accurate and precise, but it may be one or the other, or neither. Accurate, but not precise. Reached the target, but could not reproduce the result. Precise, but not accurate. Did not reach the target, but result was reproduced. Accurate and precise. Reached the target and the data was reproduced.

Measuring Accuracy Accuracy is calculated by percentage error (%E) We take the absolute value because you can’t have negative error. GP 4

Measuring Precision: Significant Figures Precision is indicated by the number of significant figures. Significant figures are those digits required to convey a result. There are two types of numbers: exact and inexact Exact numbers have defined values and possess an infinite number of significant figures because there is no limit of confidence: * There are exactly 12 eggs in a dozen * There are exactly 24 hours in a day * There are exactly 1000 grams in a kilogram Inexact number are obtained from measurement. Any number that is measured has error because: Limitations in equipment Human error

Measuring Precision: Significant Figures Example: Some laboratory balances are precise to the nearest cg (.01g). This is the limit of confidence. The measured mass shown in the figure is 335.49 g. The value 335.49 has 5 significant figures, with the hundredths place (9) being the uncertain digit. Thus, the (9) is estimated, while the other numbers are known. It would properly reported as 335.49±.01g - The actual mass could be anywhere between 335.485… g and 335.494… g. The balance is limited to two decimal places, so it rounds up or down. We use ± to include all possibilities.

Determining the Number of Significant Figures In a Result All non-zeros and zeros between non-zeros are significant 457 (3) ; 2.5 (2) ; 101 (3) ; 1005 (4) Zeros at the beginning of a number aren’t significant. They only serve to position the decimal. .02 (1) ; .00003 (1) ; 0.00001004 (4) For any number with a decimal, zeros to the right of the decimal are significant 2.200 (4) ; 3.0 (2)

Determining the Number of Significant Figures In a Result Zeros at the end of an integer may or may not be significant 130 (2 or 3), 1000 (1, 2, 3, or 4) This is based on scientific notation 130 can be written as: 1.3 x 102  2 sig figs 1.30 x 102  3 sig figs If we convert 1000 to scientific notation, it can be written as: 1 x 103  1 sig fig 1.0 x 103  2 sig figs 1.00 x 103  3 sig figs 1.000 x 103  4 sig figs *Numbers that must be treated as significant CAN NOT disappear in scientific notation

Calculations Involving Significant Figures You can not get exact results using inexact numbers Multiplication and division Result can only have as many significant figures as the least precise number 6.2251 𝑐𝑚 𝑥 𝟓.𝟖𝟐 𝑐𝑚=36.230082 𝑐 𝑚 2 =36.2 𝑐 𝑚 2 (3 s.f.) 105.86643 𝑚 0.𝟗𝟖 𝑠 =108.0269694 𝑚 𝑠 =110 𝑚 𝑠 𝑜𝑟 1.1 𝑥 1 0 2 𝑚 𝑠 (2 s.f.) 43270.0 𝑘𝑔 𝑥 𝟒 𝑚 𝑠 2 =173080 𝑘𝑔 𝑚 𝑠 2 =200000 𝑘𝑔 𝑚 𝑠 2 𝑜𝑟 2 𝑥 1 0 5 𝑘𝑔 𝑚 𝑠 2 (1 s.f.)

Calculations Involving Significant Figures Addition and Subtraction Result must have as many digits to the right of the decimal as the least precise number 20.4 1.322 83 + 104.722 211.942 212 GP 5

Limit of certainty is the ones place Mixed Operations H=10.000 cm W = .40 cm L = 31.00 cm Volume of rectangle ? V = LWH = 124 cm3 ≈ 120 cm3 or 1.2 x 102 cm3 Surface area (SA = 2WH + 2LH + 2LW) ? note: constants in an equation are exact numbers =2 4.0𝑐 𝑚 2 +2 310.0𝑐 𝑚 2 +2(1𝟐.4𝑐 𝑚 2 ) Limit of certainty is the ones place =8.0𝑐 𝑚 2 +620.0𝑐 𝑚 2 +2𝟒.8𝑐 𝑚 2 =65𝟐.8𝑐 𝑚 2 =653𝑐 𝑚 2

Dimensional Analysis Dimensional analysis is an algebraic method used to convert between different units Conversion factors are required Conversion factors are exact numbers which are equalities between one unit set and another. For example, we can convert between inches and feet. The conversion factor can be written as: In other words, there are 12 inches per 1 foot, or 1 foot per 12 inches.

Dimensional Analysis conversion factor (s) Example. How many feet are there in 56 inches? Our given unit of length is inches Our desired unit of length is feet We will use a conversion factor that equates inches and feet to obtain units of feet. The conversion factor must be arranged such that the desired units are ‘on top’ 𝟓𝟔 𝑖𝑛𝑐ℎ𝑒𝑠 𝑥 1 𝑓𝑜𝑜𝑡 12 𝑖𝑛𝑐ℎ𝑒𝑠 =4.6666 𝑓𝑡 4.7 ft GP 6

High Order Exponent Unit Conversion (e.g. Cubic Units) As we previously learned, the units of volume can be expressed as cubic lengths, or as capacities. When converting between the two, it may be necessary to cube the conversion factor Ex. How many mL of water can be contained in a cubic container that is 1 m3 3 Must use this equivalence to convert between cubic length to capacity 1 𝑚 3 𝑥 𝑐𝑚 1 0 −2 𝑚 𝑥 𝒎𝑳 𝒄 𝒎 𝟑 Cube this conversion factor =1 𝑚 3 𝑥 𝒄𝒎 𝟑 𝟏 𝟎 −𝟔 𝒎 𝟑 𝑥 𝑚𝐿 𝑐 𝑚 3 =𝟏 𝒙 𝟏 𝟎 𝟔 𝒎𝑳 GP 7