The Basics Matter and Measurements

The Basics: Matter & Measurements The Scientific Method The scientific method is a process that combines observation, hypothesis, and experimentation. • Hypotheses are the initial explanation for some observed fact or facts, are based on observation and evidence, and are to be tested and revised through experimentation. • Theories are broader in scope and explain and predict many different observations that are linked by the same underlying phenomena. Theories are generally widely accepted by scientists in the field as valid explanations of phenomena until shown otherwise. • All theories are falsifiable. Reference: Interactive General Chemistry Atoms • Scientific observations that are always true are referred to as scientific laws. • Laws are statements or mathematical relationships of a phenomenon backed by experimental evidence.. What is Matter? • Matter is anything that occupies space (i.e., has volume) and has mass. • Mass is about amount of the substance and/or object. • Weight is determined by the gravitational force exerts on the object. • Law of the conservation of matter” • Matter can neither be created or destroyed. • Matter can go from one phase to another via a physical change or chemical change. • Matter has both physical and chemical properties. The Phases of Matter • SOLIDS: have a rigid shape & fixed volume. Their external shape is reflective of their atomic and molecular arrangement. • LIQUIDS: do not have a fixed shape – their shape is the container they occupy & and they have fluidity. • GASES: expand to fill their container - their volume is determined by the container they occupy & have randomly movement. • PLASMA: defined a gaseous state of matter that contains an appreciable amount of electrically charged particles & found in very high temperature environments. • Examples: naturally: stars, lightning & man-made: Television screens. Reference credit: www.baltcnet-plasmatec.org • Matter can be classified according to its: • Composition (i.e., homogenous mixture, heterogeneous mixture or pure substance) and/or Reference credit: galleyhip.com • Unit (i.e., atom, molecule) and/or • Bonding type (i.e., metallic, molecular, ionic or covalent compounds). Mixtures • Mixtures are made up of more than one substance. • Components of mixtures can be of the same phase (i.e., solid, liquid or gas) • Mixtures can be separated into their components by using physical methods based on their components’ physical properties. • Example: • Distillation takes advantage of differences in boiling points. • Types of mixtures: • Homogenous mixtures have the same composition throughout – can not distinguish individual components. • Heterogeneous mixtures compositions varies throughout. Reference: Expii Pure Substances • Pure substances have unvarying composition. • They are two types of pure substances: • Elements • Can exist as molecules such as O2, S8 • These are referred to as molecular elements. • Most elements are monoatomic in nature, such as Cu or Fe. • Represented by symbols have one or two letters for named elements • Three-letter symbols describe elements predicted or recently discovered elements whose name has not yet be certified by IUPAC. • Only the first letter of a chemical symbol is capitalized. • Compounds (molecules) • Made up of two or more elements. • Elements can be the same such as P4 or different such as CO2. • Compounds can be classified as molecular, ionic or metallic. Reference credit:pic2fly.com Pure Substances: Groupings Based on Bond Type Molecular compounds: • compounds whose particles are molecules made of only nonmetals.. • Bonding type is covalent (i.e., electrons are shared between atoms to make a bond.) • Examples: CO2 (carbon dioxide) & NH3 (ammonia) Reference credit: 2006 encyclopedia Britannica, Inc. Ionic compounds: • compounds whose particles are cations (metals) and anions (nonmetals). • Examples: CaCl2 (calcium chloride) & CuBr (copper (I) bromide) Metallic compounds: • Compounds whose particles are composed of metal atoms or cations. (i.e., alloys) are molecules composed of metal elements. • Bonding type is metallic (i.e., sharing of electrons by metals throughout the compound). • Examples: Brass, Steel & Gold (Au) Reference credit: online.science.psu.edu Reference credit: avjky.tripod.com Matter & Its Properties • Matter can go from one phase to another via a physical change or chemical change. • Matter has both physical and chemical properties. • A physical change is a phase change. • No new compound or element is formed. • Example: dry ice (solid CO2) subliming to a carbon dioxide gas (gaseous CO2). • A chemical change is when new substances are formed in a chemical reaction. • There is a change in physical and chemical properties • Example: Wood burning. Intensive vs. Extensive Physical Properties of Matter • Physical Properties of matter can be classified as being either extensive in nature or intensive in nature. • Extensive physical properties are DEPENENT on the amount of the substances. • Example: mass, volume, color intensity, enthalpy (i.e. heat). • Intensive physical properties are INDEPENDENT of the amount of the substance. • Example: density, boiling point, melting point, color voltage (i.e. electrical potential). Physical Properties Chemical Properties • Density • Acidity (pH) • Color • Flammability • Particle size • Reactivity • Melting point • Toxicity • Boiling point • Oxidation Temperature is not Heat • Temperature is defined as the measure of hotness or coldness of an object (degrees Celsius, or degrees Fahrenheit, or kelvins). • Whereas heat is a form of energy that flows between two objects of different temperatures. • Energy flows from HOT to COLD • Directionality of heat transfer (2nd Law of Thermodynamics) • Temperature is result of the movement of matter • Is related to the average translational kinetic energy per molecule in a substance. • Cold: slow movement of particles • Hot: fast movement of particle Reference credit: ces.fau.edu Temperature: Units • Temperature is an intensive physical property of matter that measure coldness or hotness of an object or substance. • There are several different temperature scales: • Fahrenheit (oF) • Celsius (oC) • Kelvin (K) • SI unit of temperature • Kelvin = 273.15 + oC • Example: Water boils at: • 212 oF • 100 oC • 373.15 K Reference credit: OpenStax2017 Science Measurements 1. All measurements consist of two parts: a. Scalar or unit b. Numerical value Example: 25.0 cm or 1.00 ft. 2. Numerical values reflect: a. the precision of the instrument used to make the measurement or b. the type of data and measurements collected qualitative: observations quantitative: precision & accuracy Accuracy & Precision • Accuracy of a measurement is determined by how close the measure is to the accepted value. • Precision of a measurement is how similar the measures values are to it each other. • Precision of measurements is dependent on the device (i.e., glassware, tape measure, etc.) used to make the instrument. Reference credit: OpenStax2017 Metric System Gram (g) Liter (L) Meter (m) Giga Mega kilo Hecta Deca G M k H D | | | | | 109 106 103 102 101 It takes 10 g = 1 Dg It takes 1000 m = 1 km It takes 1 x 109 L = 1 GL | | deci centi milli micro nano d c m  n | | | | | 101 102 103 106 109 It takes 10 dg = 1 g It takes 1000 mm = 1 It takes 1 x 109 nL = 1 L SI Prefixes Prefix Abbreviation 12 Milli- m 10 −3 10 9 Micro- μ 10 −6 M 10 6 Nano- n 10 −9 Kilo- k 10 3 Pico- p 10 Deci- d 10 −1 Femto- f 10 –15 Centi- c 10 −2 Atto- a 10 –18 Prefix Abbreviation Meaning Tera- T 10 Giga- G Mega- Reference: Interactive General Chemistry Atoms Meaning −12 Significant Figures: Overview • Significant values always: • Nonzero integers • Captive zeros • Trailing zeros: • when to the right of a decimal point • When written in scientific notation, • Not significant values: • Leading zeroes • Trailing zeros • When they follow a nonzero integer to the left of the decimal point. Reference credit: OpenStax2017 Significant Figures & Mathematical Operations • Multiplication & Division Examples: 3.56 cm x 4 cm = 14.24 cm2 but you would report 10 cm2 2.050 mL x 5.0 mL = 10.250 mL2 but you would report 1.0 x 101 mL2 using scientific notation removes the ambiguity 45.0 cm / 9 cm = 5.0 cm but you would report 5 cm (4.5 ft) x (12 in/1 ft) x (2.54 cm/1 in) x (1 m/100cm) = 1.3716 m but you would report 1.4 m Significant Figures & Mathematical Operations • Addition & Subtractions Examples: 2 3 . 4 6 7 in + 3 1 3 . 2 1 in 3 3 6 . 6 7 7 in but you would report 336.68 in 457 cm 0 . 6 8 cm 4 5 6 . 3 2 cm but you would report 456 cm Conversion Factors & Dimensional Analysis • Conversion Factors • Is a ratio of two equivalent quantities that have different measurement units (scalars). • Examples: • • • • 100 cm = 1 m 2.54 cm = 1 in 12 in = 1 ft I mol = 6.02 x 1023 items • Dimensional Analysis (Factor Label) • Problem solving strategy using conversion factors. DENSITY: Physical Property of Matter Density = 𝐦𝐚𝐬𝐬 𝐯𝐨𝐥𝐮𝐦𝐞 •An INTENSIVE physical property. •The physical property does not depend on amount of substance. •The physical properties of mass & volume that determine a substance’s density are EXTENSIVE. •Extensive physical properties are dependent on amount. • Densities of liquid & gases are affected by temperature. Problem: The density of octane (C8H18) is 7.00 lb/gal a) what is density in mg/cm3? b) what is the mass in grams of 1.25 liters of octane? Problem: An iceberg has a volume of 7655 cubic feet. If the density of ice is 0.9985 g/mL what is the mass in kilograms of the iceberg?