Chapter 1 Chemistry Notes Fall 2024 PDF

Introduction: Matter, Energy and Measurement Chapter 1 Outline of Chapter 1 What is matter? 1.1 The Study of Chemistry 1.2 Classifications of Matter How do we distinguish between the properties of matter? 1.3 Properties of Matter...

Introduction: Matter, Energy and Measurement Chapter 1 Outline of Chapter 1 What is matter? 1.1 The Study of Chemistry 1.2 Classifications of Matter How do we distinguish between the properties of matter? 1.3 Properties of Matter 1.4 The Nature of Energy (Skip this section…covered in Ch. 5) What measurements are used in scientific calculations? 1.5 Units of Measurement How do we use measurements in scientific calculations? 1.6 Uncertainty in Measurements 1.7 Dimensional Analysis What is matter? Sections 1.1 – 1.2 The Study of Chemistry The Atomic and Molecular Perspective Chemistry is the study of matter, its properties, and the changes it undergoes. What is Matter? Matter is anything that has mass and occupies space. All matter is made up of substances called elements, which have specific chemical and physical properties and cannot be broken down into other substances through ordinary chemical reactions. The Study of Chemistry The Atomic and Molecular Perspective Chemistry explains properties and Non-Chemist behavior of matter based on structure and events at both the: atomic AND Chemist molecular levels Example: ice melting 𝒉𝒆𝒂𝒕 H2O (s) H2O (l) The Study of Chemistry The Atomic and Molecular Perspective NaCl (table salt) Atoms are the building blocks of matter (Ex: H, O, C, Na). Each element is made of the same kind of atom (Ex: C, O2). A molecule is made of two or more atoms joined together (Ex: O2, NaCl, CO2) A compound is made of two or more different types of elements (Ex: NaCl, CO2). Classifications of Matter Pure Substances Pure substances are substances with a fixed composition and unique properties. They cannot be separated by physical or mechanical means (grinding, fingers, filtering, magnets, Elements and compounds distillation, etc.). (& molecules) are pure substances. Classifications of Matter monatomic diatomic Elements Elements, if not mixed with other elements or compounds, are considered pure substances and are composed of only one type of atom (but can have more than one of that atom). Examples: iron metal (Fe) oxygen gas (O2) liquid bromine (Br2) Classifications of Matter Elements For the exam, you will need to know names and symbols for: #1-38, 40, 42, 46-48, 50, 51, 53-56, 60, 74,78-80, 82, 83, 86, 92, 94 *See Handout on Canvas (a Periodic table will be provided on exam).  Concept Check Element Symbols Match each chemical symbol with its name. Hint: Each symbol will only have one name. C Cl Possible Answer Choices: Ca cadmium chlorine Co calcium chromium carbon cobalt cesium copper Classifications of Matter States of Matter Matter can exist in three states: (1) solid (s) (2) liquid (l) (3) gas (g) Any substance can exist in the solid, liquid, or gas state. Depends on: temperature Try to envision the motion of pressure atoms/molecules in solids, liquids, and gases. How do they differ? Classifications of Matter States of Matter Physical States at Room Temperature What happens when temperature is (25°C or 298 K) INCREASED? Classifications of Matter Compounds Compounds are combinations of atoms of more than one type of element. Compounds, if not mixed with other compounds or elements, are also considered pure substances. Examples: sodium chloride (NaCl) ammonia (NH3) Classifications of Matter Compounds Law of Constant Composition (Definite Proportions) states the elemental composition of a compound is always the same. Example: Pure water from any source (nature or laboratory) is composed of: 1. two hydrogen (H) atoms 2. one oxygen (O) atom Classifications of Matter Compounds Compounds are combinations of atoms of more than one type of element. All compounds are molecules! Molecules are group of atoms (two or more) bonded chemically. Not all molecules are compounds! H2 (molecule only) NH3 (compound & molecule) O2 (molecule only) CH4 (compound & molecule)  Concept Check Elements and Compounds Classify the substances below as either an element or a compound. Hint: Consider the chemical formula of each substance below. Not all answer choices will be used. Possible Answer Element Compound Choices: air ammonia hydrogen gas iron metal granite sodium chloride Announcements Upcoming Register for MC!! CHM 114 Honors students will receive an email back soon. ✓ Lecture: Chapters 1 & 2 (possibly 3) ✓ Recitation: Act. 1 DUE 08/28 at 11:59 pm MST: ✓ Chapter 1 CC ✓ Lab: Lab 1 ✓ Chapter 1 HW ✓ Dr. Chandrakanthan ([email protected]) DUE 09/03 at 11:59 pm MST: In-Person (PSD 304F)/Zoom Office ✓ Chapter 2 CC Hours: ✓ Chapter 2 HW T Th 10:30 am – 11:30 am MST T Th 1:30 pm – 2:30 pm MST 09/02: Labor Day Observed ✓ Course Slack Page: (link soon to be announced) Online Unit 1 Exam is on Thursday, September 19th from 6:00 am – 11:59 pm MST in Canvas. Copyright © 2020 Arizona Board of Regents. Classifications of Matter Mixtures Mixtures are comprised two or more pure substances that exist together but are not combined chemically. Components of a mixture (homogeneous or heterogeneous) can be separated by physical or mechanical means. Properties of Matter Separation of Mixtures Components of a mixture CAN be separated by physical processes.These processes take advantage of the differences in physical properties of the gold iron components. Magnetic Properties: mixture of iron and sulfur Solubility: iron dissolves in HCl but gold does not Phase: use filtration to separate a solid phase from a liquid phase. Properties of Matter Separation of Mixtures Ability to form gases (differences in boiling point): distillation Examples: desalination of ocean water, making alcoholic beverages, separation of fuel from crude oil Ability to adhere to a surface or difference in solubility in a solvent: chromatography Examples: water softener, blood testing, forensic testing (blood, cloth, arson) Classifications of Matter Types of Mixtures Homogeneous Mixtures (also called solutions) Constant, uniform composition throughout Percent of each component can vary Single phase Examples: air, salt water, metal alloys, bronze, gasoline, vinegar Heterogeneous Mixtures Non-uniform composition May contain multiple phases Examples: granite, potting soil, oil & water, iron & sulfur, ice cubes in water Classifications of Matter Mixtures  Concept Check Classification of Matter Below is a sample of Roman concrete. Which classification of matter best describes this sample? A. element B. compound C. homogeneous mixture D. heterogeneous mixture How do we distinguish between the properties of matter? Section 1.3 Properties of Matter Physical properties can be Chemical properties can only observed without changing a be observed when a substance is substance into another substance. changed into another substance. Examples: boiling point, density, Examples: Flammability, mass, volume corrosiveness, reactivity with acid Properties of Matter Intensive (Bulk) Properties Extensive Properties Are independent of the amount Are dependent upon the amount of the substance that is present of the substance present. (useful for identifying Examples: mass, volume, substances). internal energy, heat capacity Examples: density, boiling/melting point, color, hardness, refractive index, temperature Properties of Matter Physical and Chemical Changes Physical Changes Chemical Changes These are changes in matter Also known as a chemical that do not change the reaction composition of a substance but Chemical changes result in new do change appearance or substances. physical state. Combustion, oxidation, Changes of state, temperature, decomposition, etc. volume, etc. Example: Example: 2 H2O (g) → 2 H2 (g) + O2 (g) H2O (l) → H2O (g) Properties of Matter Physical and Chemical Changes Physical Changes Chemical Changes At 25oC, dry ice sublimes, ice melts. Converting elements to compounds CO2 (s) → CO2 (g) 2 Na (s) + Cl2 (g) → 2 NaCl (s) H2O (s) → H2O (l) Properties of Matter Review of Phase Changes  Concept Check Physical and Chemical Changes Classify the following changes as chemical or physical. Aluminum is hammered into a sheet Important: Mixtures Copper is oxidized by concentrated nitric acid can have properties distinct from any of the Metabolism of ethanol individual substances Dry ice sublimating at room temperature making up the mixture. Dissolving sodium chloride in water Hydrogen gas burning in the presence of oxygen Liquid nitrogen boiling at room temperature Rust forming on an iron pipe What measurements are used in scientific calculations? Section 1.5 Units of Measurement Measurements are used to communicate quantitative information. Metric system is used for scientific measurements. A measurement consists of: (1) a number (2) an appropriate unit Examples: 125C, 298.15 K 10.0 kg, 10.0 mg 1.5 nm, 1.004 m 0.936 g/mL, 10 cm3 Units of Measurement SI Units Système International d’Unités (particular choice of metric units for use in scientific measurements)—A different base unit is used for each quantity. Know for Exam! Other units can be expressed as some combination of these 7 base units (derived units). Units of Measurement SI Units What is the size of Earth? Often, exponential (scientific) notation to express large and small numbers conveniently. 1.27 x 107 m Of an atom of lithium? 1.52 x 10─10 m Units of Measurement SI Units Metric prefixes are based on powers of 10 and can be used in Know for Exam! conjunction with or in lieu of exponential notation. Example: 1 x 103 g 1 kg  Concept Check Metric Prefixes Convert the measurements below. Part A: 5 g to mg Part B: 650 μL to L Units of Measurement Temperature By definition, temperature (T) is a measure of the average kinetic energy of the particles in a sample. Temperature is a physical property that determines heat flow—heat always flows spontaneously from a substance at a higher temperature to one at a lower temperature. Common units: °C, K, °F (NOT using the Fahrenheit scale in this course) The Celsius scale is based on the properties of water. The Kelvin is the SI unit of temperature and is based on the properties of gases. Units of Measurement Temperature In scientific measurements, the Celsius and Kelvin scales are most often used. The Celsius scale is based on the properties of water. − 0C is the freezing point of water. − 100C is the boiling point of water. There are no negative Kelvin temperatures! − Absolute zero: The lowest possible temperature in K is 0 K (-273.15C) To convert from Celsius to Kelvin, use the equation: 𝑲 = ℃ + 𝟐𝟕𝟑. 𝟏𝟓 Units of Measurement Temperature Common Temperatures you should know in °C and K: Normal Freezing Point of Water: 0°C, 273 K Know for Exam! Normal Boiling Point of Water: 100°C, 373 K Lowest Possible Temperature: − 273.15°C, 0 K A kelvin is the Room Temperature: same size as a 25°C, 298.15 K degree Celsius.  Concept Check Temperature Liquid nitrogen, N2 (l), boils at -195.8°C. What is the boiling point of liquid nitrogen in Kelvin? Express your answer to one place past the decimal. If your answer is negative, include the sign. How do we use measurements in scientific calculations? Sections 1.5 – 1.7 Uncertainty in Measurement Exact Numbers Inexact Numbers Values which are known exactly, Any number obtained by or defined values. measurement, uncertainties ̶ 12 eggs = 1 dozen eggs always exist in measured ̶ 2.54 cm = 1 inch quantities. ̶ 1 kg = 1000 g ̶ 1 m = 100 cm ̶ Exact counts of objects Uncertainty in Measurement Significant Figures The term significant figures refers to digits that were measured. The number of significant digits describes the exactness of the measurement. Uncertainty in Measurement Significant Figures Apply these rules to determine the number of significant figures: 1. All nonzero digits are significant. Ex: 134 2. Zeroes between two significant figures are themselves significant. Ex: 2005 3. Zeroes at the beginning of a number are never significant. Ex: 005 or 0.07 4. Zeroes at the end of a number are significant if a decimal point is written in the number. Ex: 2.200 When adding/subtracting, the final answer is expressed with the same number of decimal places as the measurement with the fewest decimal places. When multiplying/dividing, the final answer is expressed with the same number of significant figures as the measurement with the fewest significant figures.  Concept Check Significant Figures Classify each measurement below according to the number of significant figures it contains. When reading a measurement number, to determine the number of significant figures, read the number from left-to-right starting with the first non-zero digit. Measurement # of Sig Figs Measurement # of Sig Figs 17 L 0.80 g/L 903 K 1.76 x 108 C/g 100 cm 6.022 x 1023 atoms 0.0005610 g 10. moles Dimensional Analysis We use dimensional analysis to convert one quantity to another. Most commonly, dimensional analysis utilizes conversion factors. Conversion factors are often exact numbers and are not measurements (infinite sig figs) Examples: 1 cm3 = 1 mL 12 in = 1 ft 1000 mm = 1 m 60 sec = 1 min 1 lb = 453.592 grams 1 in = 2.54 cm Dimensional Analysis Write a conversion factor as a ratio. The conversion factor 1 in = 2.54 cm can be written two ways: Use the form of the conversion factor Unit conversions may involve one or that puts the desired unit in the more conversion factors to end up numerator: with the desired unit: Dimensional Analysis Example: Convert 8.00 meters to inches. Consider the path you would need to take to go from one unit to another. Step 1: Convert m to cm 10-2 m = 1 cm 1 cm 1 in. 8.00 m   = 315 ___in in 10-2 m 2.54 cm Step 2: Convert cm to in 1 in = 2.54 cm  Concept Check Dimensional Analysis The average speed of a nitrogen molecule (N2) in air at 25°C is 515 m/s. What is the speed of the nitrogen molecule (in km/hr)? Hint: 1 hr = 60 min; 1 min = 60 s What conversion factors will we need to use? Units of Measurement Derived SI Units: Volume and Density Density (d) is a physical property of a substance. 𝒎 m = mass 𝒅= 𝑽 V = volume It has units that are derived from the units for mass and volume (g/mL, g/cm3, kg/m3)—derived units. Example: Which has a greater density, lead or aluminum? 11.34 g/cm3 2.70 g/cm3 Units of Measurement Derived SI Units: Volume and Density Density is a physical property of a substance. 𝒎 𝒅= 𝑽 Example: If each side of the lead cube is 2.15 cm, what is the mass of the cube (in g)? **Equation Method VPb cube = length x length x length = (2.15 cm)3 = 9.94 cm3 𝒎 𝒅= Rearrange for m. 𝑽 m = dPb x VPb = (11.34 g/cm3) x (9.94 cm3) = 113 g Units of Measurement Derived SI Units: Volume and Density Density is a physical property of a substance. 𝒎 𝒅= 𝑽 Example: If each side of the lead cube is 2.15 cm, what is the mass of the cube (in g)? **Conversion Factor Method VPb cube = length x length x length = (2.15 cm)3 = 9.94 cm3 9.94 𝑐𝑚3 𝑃𝑏 11.34 𝑔 𝑃𝑏 × = 113 g 1 1 𝑐𝑚3 𝑃𝑏  Concept Check Density Calculate the volume (in mL) of 65.0 g of liquid methanol if its density is 0.791 g/mL. END OF CHAPTER 1 SLIDES