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Chapter 1

by: Brianda Hickey

Chapter 1 CHEM-UA 125

Brianda Hickey
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An outline of Chapter 1 in Chemistry: The Molecular Nature of Matter and Change, 7th Ed., by Martin S. Silberberg
General Chemistry I
Dr. Malgorzata (Margaret) Mandziuk
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This 7 page Class Notes was uploaded by Brianda Hickey on Wednesday February 3, 2016. The Class Notes belongs to CHEM-UA 125 at NYU School of Medicine taught by Dr. Malgorzata (Margaret) Mandziuk in Spring 2016. Since its upload, it has received 213 views. For similar materials see General Chemistry I in Chemistry at NYU School of Medicine.

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Date Created: 02/03/16
Chapter 1 - Keys To The Study Of Chemistry *To find more information, reference the textbook. The headings used below correspond with those in the textbook. 1.1 Chemistry: The study of matter and its properties, the changes that matter undergoes, and the energy associated with those changes. Matter: The “stuff” of the universe. Anything with mass and Volume ex. air, glass, plants Composition: The types and amounts of simpler substances that make up matter Substances: a type of matter that has a defined, fixed composition The States of Matter 3 Physical Forms (States) Solid: Fixed Shape Liquid: Varying shape, conforms to container only to the extent of liquid’s volume Gas: varying shape, conforms to container, fills entire containers Different Atomic States Solid: Particles are close together and organized Liquid: Particles are close together and disorganized Gas: Particles are far apart and disorganized Properties of Matter & Its Change Properties: The characteristics that give each substance its unique identity To identify substance: observe physical & chemical properties Physical Change (aka No change in composition) Physical Properties: Characteristics shown by a substance without being changed or interacting with another substance melting point, electrical activity, density Physical Change: When a substance alters its physical properties, NOT its composition ex. Water transitioning from solid to liquid; There is physical change but no change in composition Chemical Change (aka Yes change in composition) Chemical Properties: characteristics shown as a substance changes into or interacts with another substance flammability, corrosives etc. Chemical change (Chemical reaction): When a substance is converted into a different substance ex. Water + electrical current = hydrogen & oxygen Temperature and Changes In Matter Physical change caused by heating can generally be reversed by cooling NOT true for a chemical change ex. moisture in are + hot iron = rust. The iron cannot be returned to its original form (remove the rust) by cooling its temperature Central Theme In Chemistry We study observable (macroscopic-scale) changes in matter to understand their unobservable (atomic-scale) causes. The Importance of Energy In The Study Of Matter Energy: ability to do work Total Energy = [Potential energy] + [Kinetic Energy] Potential Energy: The energy due to the position of the object relative to other objects Kinetic Energy: The energy due to the motion of the object Energy changes form, but it is conserved Chemical Potential Energy arises from the positions and interactions of a substance’s particles. When a higher energy (less stable) substance is converted into a more stable (lower energy) substance, some potential energy is converted into Kinetic Energy 1.4 Measurement & Chemical Problem Solving SI System used in all of chemistry so scientists around the globe can communicate 7 fundamental units (Base Units) Derived units - Combinations of the seven base units ex. Derived unit for speed = (m/s). Base unit for length (m) divided by the base unite for time (s). Some Important SI Units in Chemistry Length meter (m) (V) Volume cubic meter (m^3) [SI Unit] Liter (L) & Millimeter (mL) [ Chemistry non-SI Unit] Mass Kilogram (kg) [SI Unit] Mass Weight Weight Constant Depends on local gravitational field Quality of matter cannot change acting on the object varies with altitude Time Second (s) [SI Unit] Units and Conversion Factors in Calculations All measured qualities consists of a number and a unit 5 Feet, 10 Inches… NOT 5,10 Units can be multiplied, divided, and canceled Constructing a Conversion Factor Conversion Factor: ratios used to express a quantity in different units ex. miles -> feet Use Equivalent Quantities -> 1mi = 5280 ft 1mi/5280ft or 5280ft/1mi Distance (ft) = 150 mi x 5280ft/1mi = 792,00 ft Density: a Combination of Units as a Conversion Factor Density (d) is a characteristic physical property of a substance and is the ratio of its mass to its volume mass = volume x density SI Unit = (kg/m^3) Chemistry Unit = (g/L) or (g/mL) Densities of gases are lower than densities of liquid or solids Temperature Temperature (T): The Measure of how hot or cold one object is relative to another Heat: The energy that flows from the object with higher temperature to the lower temperature 3 Types of Temperature Kelvin (K) [SI Base Unit] Freezing: 273 K Boiling 373.15 K Celsius (C) Freezing: 0 C Boiling: 100 C Fahrenheit (F) Freezing: 32 F Boiling: 212 F Extensive & Intensive Properties Extensive Properties: Dependent on the amount of space present mass & volume Intensive Properties: Independent of the amount of substance Density Example: Heat (extensive), Temperature (intensive) boiling water has more heat (energy) than a cup of boiling water, but both have same temperature 1.5 Uncertainty in Measurement: Significant Figures Every measurement includes uncertainty use (+ -) to display the uncertainty in measurement The greater the amount of significant figures, the greater is the certainty of measurement Determining What Units are Significant All numbers are significant except series used only to position the decimal point If there is a decimal point and zero 1.1300g = 5 Significant figures 6500. = 4 Sig. Fig. Only zeroes after or before decimal counts If there is NO decimal point, but there is a zero assume zeros are not significant, unless exponential notion clarifies 5300 L = 2 Sig. Fig. 5.300 X 10^3L = 4 Sig. Fig. 5.30 X 10^3L = 3 Sig. Fig. 5.3 X 10^3L = 2 Sig. Fig. terminal decimal point indicates zeros are significant 500 mL = 1 Sig. Fig 500. mL = 3 Sig. Fig. Significant Figures: Calculations and Rounding Off In calculations - keep track of Significant Figures If have too many numbers, round off Choose number of Significant Figure from the measurement with the least amount of significant figures ex. Find the Density: 3.8056 g/ 2.5 mL = 1.522 g/mL = 1.5 g/mL Rules For Arithmetic Operations 1. Multiplication and Division Answer contains same number of significant figures as there are in the measurement with fewest significant figures ex. Volume (cm^3) = 9.2cm X 6.8cm X 0.3744cm = 23.4225cm^3 = 23cm^3 2. Addition and Subtraction The answer has same number of decimal places as there are in the measurement with the fewest decimal places ex. 83.5mL + 23.28mL = 106.78mL = 106.8mL Rules of Rounding Off 1. If digit removed is more than 5 -> preceding number goes up by one 5.379 -> 5.38 ( 3 sig fig) 5.4 (s sig fig) 2. If digit removed is less than 5, preceding number remains the same 0.2413 -> 0.241 (3 sig fig) 0.2413 -> 0.24 (2 sig fig) 3. If digit removed is 5, the preceding number goes up, if odd. 17.75 -> 17.8 the preceding number remains the same if even 17.65 -> 17.6 If 5 is followed by only zeroes - Follow rule 3 17.6500 -> 17.6 If 5 followed by nonzero -> rule 1 followed 17.6513 -> 17.7 4. Always carry one or two additional sig fig through a multistep calculation and round off the final answer ONLY Precision, Accuracy, Instrumental Calibration Precision (reproductabilit); how close the measurements in a series are to each other Accuracy: how close each measurement is to the actual value Systematic Error: values that are either all higher or all lower than the actual value Random error: the absence of systematic error, produces values that are higher and lower than the actual value Calibration: comparing the measuring device with a known standard systematic error (caused by faculty equipment) may be compensated with calibration


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