CHM2045 General Chemistry I Chapter 1 PDF

CHM2045: General Chemistry I Chapter 1: Matter, Measurement, and Problem Solving Omega Nebula is ~15 light years across or 82000000000000 (8.2 × 1013) miles across. CHM2045: General Chemistry I Chapter 1: Matter, Measurement, and Problem S...

CHM2045: General Chemistry I Chapter 1: Matter, Measurement, and Problem Solving Omega Nebula is ~15 light years across or 82000000000000 (8.2 × 1013) miles across. CHM2045: General Chemistry I Chapter 1: Matter, Measurement, and Problem Solving Learning goals and key skills: Understand what constitutes matter Describe the scientific approach to knowledge Apply the two methods for classifying matter Understand what constitutes a homogenous mixture, heterogeneous mixture, pure substance, compound, and element Describe chemical and physical properties Describe chemical and physical changes Describe intensive and extensive properties Discuss energy as a fundamental part of physical and chemical changes Identify the units of measure for physical phenomena Memorize the metric system units Understand the differences between a derived unit and a base unit Describe the differences between an exact number and inexact numbers Determine the number of significant figures in a measurement and calculation Utilize dimensional analysis to solve calculations Skip Section 1.9: Analyzing and Interpreting Data 1 Atoms and Molecules Chemistry is the science that seeks to understand the behavior of matter by studying the behavior of atoms and molecules. Atoms are submicroscopic particles that are the building blocks of matter. Molecules are atoms that are held together in a specific shape. Molecules are combinations of atoms held together in a specific shape Changes in structure and composition affect properties Ethanol Ethylene Glycol Stoichiometry = C2H6O Stoichiometry = C2H6O2 Melting Point = -115 ºC Melting Point = -16 ºC Boiling Point = 78 ºC Boiling Point = 197 ºC Density = 0.79 g/cm3 Density = 1.11 g/cm3 Consumption = Intoxicating Consumption = Toxic 2 The Scientific Approach to Knowledge Scientific knowledge is based on observation and experiment Hypothesis is a tentative explanation of observations Law describes how nature behaves Theory is a model for the way nature behaves The Classification of Matter Matter is anything that occupies space and has mass. Matter is classified according to its: State (its physical form): solid, liquid, and gas Composition (the basic components that make it up): element, compound, mixtures 3 The States of Mater: Solid, Liquid, and Gas SOLID - volume and shape independent of container, rigid, incompressible LIQUID - volume independent of container, no fixed shape, incompressible GAS - no fixed volume or shape, conforms to container, compressible Classifying Matter by Composition: Pure Substance Pure Substance is made up of only one component, and its composition dose not vary from one sample to another. The components can be individual atoms or groups of atoms, but any one of these substances is always the same. Two types of pure substances: An element is a pure substance that can't be broken down into simpler substances A compound is a pure substance composed of two or more elements in a fixed definite proportion. 4 Classifying Matter by Composition: Mixture A mixture is composed of two or more components that can vary from one sample to another. Two types of mixtures: A heterogenous mixture has variable composition from one region of the mixture to another. A homogenous mixture have uniform composition throughout the sample because the atoms and/or molecules that compose them are mixed. Classifying Matter by Composition: Pure Substances and Mixtures. In these images, the blue circles represent one type of element, and the red squares represent a second type of element. Classify each circle. 5 Classifying Matter by Composition: Elements Compound, and Mixtures Thought Experiment: How would you classify the following compounds? granite mercury brass sweet tea 6 Separating Mixtures Chemists often want to separate a mixture into the its components. Mixtures are separable because the different components have different physical and chemical properties. Examples include decanting, distillation, and filtration Physical and Chemical Changes Physical changes alter the state or appearance (not the composition) of matter. Chemical changes alter the composition of matter. Atoms rearrange and transform the original substance into a different substance. 7 Physical and Chemical Properties Physical properties are properties that a substance displaces without changing its composition. Examples include color, density, mass, odor, melting point, boiling point, hardness Chemical properties are properties that substance displaces only changing its composition via a chemical change. Examples include flammability and reactivity with acid. Intensive and Extensive Properties An intensive property is a characteristic that is independent of the amount of substance. Examples include melting point, temperature, odor, color, etc. An extensive property is a characteristic that is dependent on the amount of substance. Examples include mass, volume, energy, etc. 8 Practice Problems State whether each of the following is a physical or chemical change: - mercury cools to a silvery solid, - natural gas burns with a blue flame, - ammonia and copper solution turns a deep blue, - sugar crystals dissolve in water. State whether each of the following is a physical or chemical property: - tungsten has the highest melting point, - ice floats on water, - potassium metal fizzes in water, - silver metal turns black in air. Energy: A Fundamental Part of Physical and Chemical Changes Understanding physical and chemical changes requires a discussion of energy changes and energy flow Kinetic energy is the energy associated with motion Potential energy is the energy associated with position or composition. Thermal energy is the energy associated with the temperature of an object. Energy is always conserved in a physical or chemical change; it is neither created nor destroyed (law of conservation of energy). Systems with high potential energy tend to change in ways that lower their potential energy, transferring energy to the surroundings. 9 The Units of Measure Units are standard quantities used to specify measurements. The two most common unit systems are the Imperial System and Metric System. Scientists use Si Units (Système International dʼUnités), which is based on the Metric System. Metric System Imperial System 10  &*@$%^@# $@ *@^@$%#&*#!!! ! Imperial to metric conversion factors/unit equations will be provided on exams. http://xkcd.com/526/ Standard Units - 7 Base units from which all other units can be derived from these http://physics.nist.gov/cuu/Units/current.html 11 Clarifying Temperature Scales (F  32) These equations will be C  provided on the exam 1.8 K  C  273.15 Si Units Uses Prefix Multiplies to Change the Value of the by Powers of 10. These relationships need to be memorized! 12 Practice Problems State the name and symbol for the metric prefix that represents the following: 1×106, 1×103, 1×10-9, 1×10-12 What is the symbol for the metric unit microliter? A) µL B) mL C) cL D) ML E) I am not sure General Problem-Solving Strategy Questions to ask in dimensional analysis: 1. Identify the starting point (the given information). 2. Identify the end point (what you must find). 3. Devise a way to get from the starting point to the end point using what is given as well as what you already know or can look up (use a conceptual plan) In other words, Given → Conceptual Plan → Find Practice Problem: How many grams are there in 25 lbs? Potentially useful unit equation: 1 lb = 456.6 g. Practice Problem: How many liters are there in 23.5 mL? 13 Derived Units: A Combination of Other Units A derived unit is a combination of other units Examples of derived units include: speed = length/time Si units: meters/second Commonly seen as: miles/hour, kilometers/hour volume = length3 Si units: m3 1 cm3 = 1 milliliter = 1 mL 1000 mL = 1 liter (not SI) Derived Units: Density Density of a substance is the ratio its mass to its volume. For solids, the units for density are typically reported at g/cm3 For liquids, the units for density are typically reported at g/mL For gases, the units for density are typically reported at g/L Practice Problem: Calculate the mass in grams for 4.80 cm3 of table salt. The density of table salt is 2.18 g/cm3 Practice Problem: A block of copper has a mass of 143.584 g and measures 5.05 cm by 2.55 cm by 1.25 cm. What is the density of the rectangular copper block? A) 29.0 g/cm3 B) 28.4 g/cm3 C) 8.92 g/cm3 D) 11.1 g/cm3 14 The Reliability of a Measurement All measurements have some degree of uncertainty associated with them. The number of digits reported in a measurement reflects the certainty in the measurement. - More digits, more certainty; fewer digits, less certainty. Counting Significant Figures Significant figures are used communicate the precision/certainty of a measurement. - The greater number of significant figures, the greater the certainty of the measurement. Significant Figure Rules 1. All nonzero digits are significant. 2. Interior zeroes (zeroes between two nonzero digits) are significant. 3. Leading zeroes (zeroes to the left of the first nonzero digit) are not significant. They only serve to locate the decimal point. 4. Trailing zeroes (zeroes at the end of a number) are categorized as follows: Trailing zeroes after a decimal point are always significant. Trailing zeroes before a decimal point (and after a nonzero number) are always significant. Trailing zeroes before an implied decimal point are ambiguous and should be avoided by using scientific notation. 15 Counting Significant Figures Significant Figure Rules 1. All nonzero digits are significant. 2. Interior zeroes (zeroes between two nonzero digits) are significant. 3. Leading zeroes (zeroes to the left of the first nonzero digit) are not significant. They only serve to locate the decimal point. 4. Trailing zeroes (zeroes at the end of a number) are categorized as follows: Trailing zeroes after a decimal point are always significant. Trailing zeroes before a decimal point (and after a nonzero number) are always significant. Trailing zeroes before an implied decimal point are ambiguous and should be avoided by using scientific notation. Practice Problem: How many significant figures are in the following values? 3003 g 1.00 × 106 m 0.02000 s 2000 miles 2000.00 miles Uncertainty in Measurements 2 types of numbers: inexact numbers and exact numbers Inexact numbers have uncertainty - numbers from measurements are always inexact and have some level uncertainty - due to both equipment errors and human errors. - uncertainties always exist in measured quantitates - use significant figures Exact numbers do not have uncertainty. They are derived from: - counting of objects such as 3 atoms, 10 students, etc. - defined quantities such as 12 eggs in a dozen, 1 m = 100 cm, etc. - integral numbers from an equation such as "diameter = 2 × radius" 16 Practice Problems How many significant figures are in each number? 554 km 7 pennies 1.01×105 m 0.00099 s 1.45000 km 21,000 m Significant Figures in Calculations Calculations are only as precise as the least significant measurement Multiplication/division: Report to the least number of significant figures 6.221 cm × 5.2 cm = 32.3492 = 32 cm2 Addition/subtraction: Report to the fewest number of decimal places 20.4 g - 1.322 g = 19.078 = 19.1 g Rounding off: looking at the digits to be dropped: 1. If leftmost digit < 5, leave preceding number the same. 2. If leftmost digit  5, increase preceding number by 1. If multiple steps calculations are required, retain an extra significant figure until the end of the calculation to prevent rounding errors. Using a calculator, enter numbers and only round to the final answer. 17 Multiple Measurements and Uncertainty Accuracy and Precision and are used to described the uncertainty in repeated measurements. - Accuracy refers to how close the measured value is to the actual value. - Precision refers to how close a series of measurements are to one another (or how reproducible they are). Dimensional Analysis Uses Units to Guide Problem Solving This is the strategy that will be employed throughout the entire course!!! Key concept or tool is the conversion factor!!! - ratios to convert between two different types units (grams to milliliters) or the same quantity (inches to centimeters) - exact numbers so do not factor into sig figs - can written as because they are the same value and related to unity: 1 inch 2.54 cm and 2.54 cm 1 inch 18 Units Raised to a Power When building factors for units raised to a power, remember to raise both the number and the unit to the power. There are 2 ways to consider units raised to a power. single step multi step 2.54 cm = 1 in 2.54cm =1 in (2.54 cm)2 = (1 in)2 2.54cm =1 in (2.54)2 cm2 = 12 in2 6.45 cm2 = 1 in2 Practice Problem: How many in2 are in 255 cm2? Potentially useful unit equation: 2.54 cm = 1 in. Practice Problems Practice Problem: Ethyl alcohol, one of the components of beer and wine, has a density of 0.789 g/mL. Determine the volume (in mL) of a sample of ethyl alcohol with a mass of 35.5 g. A) 28.0 mL B) 2.80 mL C) 45.0 mL D) 4.50 mL E) none of the above Practice Problem: If a bronze solid has a volume of 25.5 cm3, what is the volume in cubic inches? (1 inch = 2.54 cm) A) 1.56 inch3 B) 3.95 inch3 C) 10.0 inch3 D) 64.8 inch3 E) 418 inch3 19 20