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States of Matter Chapter 3 Matter and Energy States of Matter (cont.) Gases 1 Liquids Properties • Characteristics of the substance under observation • Properties are: – Directly observable – The way something interacts with other substances in the universe Universe Classified • Matter: the part of the universe that has mass and volume • Chemistry is the study of matter – The properties of different types of matter – The wayy matter changes g and behaves when influenced by other matter and/or energy Properties of Matter • Physical Properties: the characteristics of matter that can be changed without changing its composition – Characteristics that are directly observable • Chemical Properties: the characteristics that determine how the composition of matter changes as a result of contact with other matter or the influence of energy – Characteristics that describe the behavior of matter 2 Chemical Properties Chemical Properties (cont.) • One commonly cited chemical property is flammability, the ease with which a substance burns in a flame. Burning is a chemical reaction. Classify Each of the following as a Physical or Chemical Property • Ethyl alcohol boiling at 78°C. • Hardness of a diamond. • Sugar fermenting to form ethyl alcohol. Classify Each of the following as a Physical or Chemical Property (cont.) • Ethyl alcohol boiling at 78°C. – Physical property: boiling point is a associated with a phase change. It describes an inherent characteristic of alcohol. • Hardness of a diamond. – Physical property: describes an inherent characteristic of diamond – hardness • Sugar fermenting to form ethyl alcohol. – Chemical property: describes behavior of sugar – forming a new substance (ethyl alcohol) through a chemical reaction 3 Changes in Matter Chemical Change • Physical changes: changes to matter that do not result lt in i a change h the th fundamental f d t l components t that make up the substance – State changes: boiling, melting, condensing • Chemical changes: changes that involve a change in the fundamental components p of the substance – Produce new substances – Chemical reactions occur – Reactants → Products Chemical Change (cont.) • Chemical change involves a chemical reaction. At least one new substance is formed. Classify Each of the following as a Physical or Chemical Change • Iron metal melting. • Iron combining with oxygen to form rust. • Sugar fermenting to form ethyl alcohol. 4 Classify Each of the following as a Physical or Chemical Change (cont.) • Iron metal melting. – Physical change: describes a state change, but the material is still iron • Iron combining with oxygen to form rust. – Chemical change: describes how iron and oxygen react to make a new substance, rust • Sugar fermenting to form ethyl alcohol. – Chemical change: describes how sugar forms a new substance (ethyl alcohol) via a chemical reaction Classification of Matter Matter Pure Substance Constant Composition Homogeneous Mixture Variable Composition • Homogeneous: uniform throughout, appears to be one thing – Pure substances b – Solutions (homogeneous mixtures) • Heterogeneous: non-uniform, contains regions with different properties than other regions Elements and Compounds • Elements: substances that cannot be broken d down into i simpler i l substances b by b chemical h i l reactions • Most substances are chemical combinations of elements. These combinations are called compounds. – Compounds are made of elements – Compounds can be broken down into elements – Properties of the compound not related to the properties of the elements that compose it – Same chemical composition at all times Pure Substances • Pure substances – All samples have the same physical and chemical properties – Constant composition: all samples have the same composition – Homogeneous – Separate into components based on chemical properties 5 Mixtures Pure Substances vs. Mixtures • Mixtures – Different samples may show different properties – Variable composition – Homogeneous or heterogeneous p into components p based on p physical y – Separate properties • All mixtures are made of pure substances Solutions Solutions (cont.) • A solution is a homogeneous mixture. • Phase can be gaseous, liquid, or solid. 6 Identity Each of the following as a Pure Substance, Homogeneous, Mixture, or Heterogeneous Mixture. Identity Each of the following as a Pure Substance, Homogeneous Mixture, or Heterogeneous Mixture (cont.) • Gasoline • Gasoline – A homogenous mixture • A stream with gravel on the bottom • A stream with gravel on the bottom • Copper metal – A heterogeneous mixture • Copper metal – A pure substance (all elements are pure substances) Separation of Mixtures Separation of Mixture (cont.) • Mixtures can be separated based on different pphysical y pproperties p of the components p – Physical change Different Physical Property Technique Boiling point Distillation State of matter (solid/liquid/gas) Adherence to a surface Chromatography Volatility Evaporation Filt ti Filtration 7 Separation of a Mixture (cont.) Another Look at Matter Energy and Energy Changes Energy and Energy Changes • Energy: ability to do work or produce heat • Potential Energy: energy due to composition or position • Energy may affect matter. • Kinetic Energy: energy due to motion – - ½ mv2 – Chemical Chemical, mechanical mechanical, thermal thermal, electrical, electrical radiant, radiant sound, nuclear – Potential and kinetic – e.g. Raise its temperature, eventually causing a state g , or cause a chemical change g such as change, decomposition • All physical changes and chemical changes involve energy changes. 8 Energy and Energy Changes (cont.) Temperature and Heat • Law of Conservation of Energy: energy can be converted from one form to another, but cannot be created or destroyed • Heat: a flow of energy due to a temperature difference • Temperature: a measure of the random motions of the components of a substance Temperature and Heat (cont.) Exothermic vs. Endothermic • System: that part of the universe that we wish to study t d • Surroundings: everything else in the universe • Exothermic process: a process that results in the evolution of heat - Example: when a match is struck, it is an exothermic p process because energy gy is produced p as heat. • Endothermic process: absorbs heat - Example: melting ice to form liquid water is an endothermic process because the ice absorbs heat in order to melt 9 Exothermic Process Units of Energy • One calorie = amount of energy needed to raise the temperature of one gram of water by 1°C – kcal = energy needed to raise the temperature of 1000 g of water 1°C • joule – 4.184 J = 1 cal • In nutrition, calories are capitalized. – 1 Cal = 1 kcal Example - Converting Calories to Joules Energy & Temperature of Matter Convert 60.1 cal to joules. • The amount the temperature of an object increases depends on the amount of heat added (q). (q) 1 cal = 4.184 joules 4.184 J 60 1cal × 60.1cal = 251J 1 cal – If you double the added heat energy the temperature will increase twice as much. • The amount the temperature of an object increases when heat is added depends on its mass – Iff you ddouble bl the h mass iit will ill take k twice i as muchh heat h energy to raise the temperature the same amount. 10 Specific Heat Capacity Specific Heat Capacity • Specific heat (s): the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius By definition , the specific heat of water is 4.184 J g °C Amount of Heat = Specific Heat x Mass x Temperature Change Q = s x m x ∆T Example #1: Calculate the amount of heat energy (in joules) needed to raise the temperature of 7.40 g of water from 29.0°C to 46.0°C. Example #1 (cont.) Specific heat of water = 4.184 J g °C Mass = 7.40 g Temperature change = 46.0°C – 29.0°C = 17.0°C Q = s • m • ∆T Heat = 4.184 J × 7.40g ×17.0°C = 526 J g °C 11 Example #2 A 1.6 g sample of metal that appears to be gold ld requires i 5.8 5 8 J to raise i the h temperature from 23°C to 41°C. Is the metal pure gold? Example #2 Q = s × m × ∆T Q m × ∆T ∆T = 41°C - 23°C = 18°C s= s= J 5.8 J = 0.20 1.6 g x 18°C g °C Table 10.1 lists the specific heat of gold as 0.13 Therefore the metal cannot be pure gold. J g °C 12