Chapter 3 Matter and Energy PDF

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1/31/21 Chem 101 Chapter 3: Matter and Energy 

Classifying Matter o Matter- anything occupying space (has volume) and having mass  Some core assumptions about matter  All matter is composed from atoms  Atoms are much too small to be seen, even with the best microscopes  There is more than one type of atoms  Atoms cannot be created or destroyed nor converted to another element  Atoms can bond together o What can atoms make?  Atoms can bond together or with other types of atoms  In different ratios  Uniformly or nonuniformly o Classifying matter by its particles  Pure substances- composed of just one type of particle  Mixtures- composed of more than one type of particle  Elements- composed of particles with just one type of atom  Compounds- composed of particles that have more than one type of atom in them o Homogeneous mixtures- different particles uniformly mixed (on macroscopic level, cannot tell if there are different types) o Heterogeneous mixtures- different particles mixed nonuniformly (on macroscopic level, easily see they are different)  Elements- composed of particles that are individual atoms (most are this)  Diatomic elements- composed of two atoms of the same type of element bonded together to form a molecule (molecular elements)  Ex. Hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, iodine  Tricky examples:  Bronze- made from tin and copper together o Can be separated back into copper and tin by careful heating via liquation  Aqueous copper sulfate- made my adding copper sulfate to water and mixing

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o Can be separated back into copper sulfate and water by evaporating the water Water- made by reacting a mixture of hydrogen and oxygen gases Sodium chloride- made in a chemical reaction of chlorine gas and solid sodium metal

o Tests  Visual  Can see 2+ different components = heterozygous mixture  Appears to be just one component = element, compound, or homogeneous mixture  Elemental  Use mass spectroscopy to break up into its atoms and detect o Just one element present = element o 2 or more elements = compound or mixture o Fixed element ratio = compound (mass spectroscopy determines ratio by running many samples)  Separation  Filter, distillation, evaporation, use magnet, eddy current, etc. o If sample can be separated = mixture o Summary  Elements  Particle level- all the same type of atom, may bond in pairs  Chemical formula- contains only one chemical symbol, may have a subscript  Chemical name- a single element name, one of the 118 elements found on the periodic table  Compounds  Particle level- two or more different atoms bonded together  Chemical formula- contains two or more chemical symbols  Chemical name- often contains parts of two element names in it (many exceptions)  Mixtures  Particle level- random mix of different atoms and bonded atoms  Chemical formula- does not exist  Chemical name- does not exist o Can be separates by physical means  Has components that are just part of the substance  Varying composition (different types) Physical vs. Chemical Properties o Physical change  Change in the form of a substance, not in its chemical composition  Ex: boiling or freezing water

Can be used to separate a mixture into pure compounds, but it will not break compounds into elements  Distillation  Filtration  Chromatography  At the atomic level, a change in the in the arrangement of the tiny particles (no bonds formed or broken) Chemical change  A given substance becomes a new substance or substances with different properties and different composition  Ex: Bunsen burner (methane reacts with oxygen to form carbon dioxide and water)  At the atomic level, a change in bonds of the particles  Bonds are formed, broken, or both Classifying properties of matter  Physical properties  Anything that can be observed/measured without chemically changing the substance o Color, odor, texture o Density, melting point  Chemical properties  Anything describing a substance’s ability to form new substances o Gasoline and oxygen make smoke and release heat o Iron makes rust with air Atomic change  Using the model of matter being particles, then any change can either be  A rearrangement of the particles (physical change) o Close particles to far apart o Separate particles to mixed  A change in how they are stuck together (chemical change) o Separate particles that change to particles bonded together What we observe  Copper undergoes a physical change when it is melted, molded, cut, or stamped to make coins  Copper undergoes a chemical change when it reacts with substances in the air and forms copper verdigris What happens to the atoms  Copper undergoes a physical change when it is melted, molded, cut, or stamped to make coins  It is still the same atoms at the atomic level, just rearranged  It undergoes a chemical change when it reacts with substances in the air and forms copper verdigris 

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The atoms have bonded with other atoms to form new compounds o Guidelines for chemical change  Change is not reversible  Unpredicted color change  New odors  Gases made (bubbles) when not heating it  Solids made (becomes cloudy or see crystals) when not cooling it o Confusion  Physical properties change during chemical changes  When water boils (physical change) o Water (liquid) becomes steam (gas) o This is a physical change, because the substance has not changed, only its physical properties gave  When gasoline burns (chemical change) o Gas (liquid) becomes carbon dioxide and water (gases) o Not a physical change but the new substances have different physical properties than the original  Need to use context clues to decide if it is just a physical change, or if the substance is changing into another one and it is really a chemical change o Guidelines for physical change  The change is reversible (ends up as the original substance so it probably never changed)  A predictable color change, mixing colors together  Gases made (bubbles) because its being heated  Solids made (becomes cloudy or see crystals) because you are cooling it Conversion of Mass o Law of conservation of mass  The total mass of materials is not affected by a chemical change in those materials  Mass of A + mass of B = mass of C + mass of D o A+B=C+D  Mass of all substances before chemical change equals mass of all substances after  Write down the reaction, list of reactants yields list of products  Write the numbers of the masses you know below each substance  Use algebra to figure out the mass you don’t know States of Matter o Solid  Particles have strong attraction 

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Particles are pulled close together and held tightly, often in a repeated pattern of packing Fixed shape and fixed volume

 o Liquid  Particles have weaker attraction  Particles still held together, but able to move independently and slide past one another  Shape capable of changing  Fixed volume o Gas  Particles have weakest attraction  Particles are far apart and move fast and freely  Shape capable of changing and volume capable of changing  Compressible o State changes are physical  Particles rearrange their spacing  During state changes, both states co-exist Energy o Energy- the ability to do work  Supplying heat is one kind of work  Moving an object is another kind of work o Potential energy- stored energy due to position  In chemistry we focus on particles, so this is energy related to the location of the particles (and where they have the potential to be)  All particles of matter are weakly attracted to each other  When particles of matter are far apart, they have a lot of potential energy (they have the potential to move together, because they are attracted)  When particles of matter are close together, they have less potential energy o Kinetic energy- energy of motion  In chemistry, we focus on particles, so this is energy related to how fast the particles are moving, which is related to its temperature o Types of energy within matter  Potential energy- stored energy due to position  A ball at top of a hill (with the potential to fall)  A stretched rubber band (with the potential to snap back)  Within matter, two attracted particles with the potential of snapping back together and bonding o Relates to state: a gas substance has a higher potential energy than its solid state o Relates to chemical bonds: unbonded particles have higher potential energy than bonded particles o Relates to electrons within atoms: an atom with electrons far from the nucleus has higher potential energy than one with electrons close to the nucleus

Kinetic energy- energy of motion  It can be the whole object moving (translational)  Within matter, movement of the atoms within an object (vibrational or rotational) o Relates to the temperature of a sample: a hot sample has more kinetic energy than a cold one Units of energy  The common unit of energy in science is the Joule (J)  It takes humans 10 Joules to quickly lift a 1 kg mass to 1 meter above the floor  Another common unit is calorie (cal)  1 cal = 4.184 J  Note: nutritional calories are not equivalent to calories  1 food calorie (Cal) = 1 kcal = 1000 cal Temperature units  Temperature is related to kinetic energy  Celsius- referenced to water’s freezing and boiling o C = F – 32 (5/9)  Fahrenheit- referenced to saltwater’s freezing o F = C (9/5) + 32  Kelvin- used in chemistry because 0K corresponds to the molecule having no kinetic energy (a molecule still has plenty of KE at 0 C) o K = C + 273.15 Heat energy  Temperature does not equal heat  Temperature is a measure of the average kinetic energy belonging to a sample (per particle)  Heat energy is the method in which energy is transferred from one object to another  When heat is transferred to an object, it results in a temperature change or a phase change or a chemical change  Heat energy flows from a warmer object to a colder object until both objects reach a thermal equilibrium Law of conservation of energy  Energy cannot be created or destroyed  It can, however, be transferred to another object or type of energy  In our case o A system of a collection of matter o The surroundings of more matter  When something transfers energy into matter, it absorbs energy  KE, PE, or both must increase  Endothermic  When energy is transferred from matter, it releases energy  KE, PE, or both must decrease  Exothermic 

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o State change or temperature change  When heat is added to a sample of matter, what happens depends on the temperature of the sample  If the sample is at its melting point temp or boiling point, a state change occurs (potential energy changing)  Otherwise, a temperature change occurs (kinetic energy changing)  While a sample is undergoing a state change (melting, boiling, condensation, freezing), it does not change temperature o Energy and chemical reactions  Energy is needed to break bonds and is released when bonds are made  Each reaction has a different characteristic energy needed to make or break the bonds  We call this energy deltaHrxn (delta H of reaction or enthalpy of reaction)  ΔHrxn is positive when a reaction absorbs energy (increases) from the surroundings to occur (endothermic)  ΔHrxn is negative when a reaction releases energy (decreases) to the surroundings (exothermic) o Endothermic  When energy is absorbed, it is called an endoergic process  If it absorbs heat energy, it is endothermic  When energy is absorbed one of three things must happen for energy to be conserved:  The temperature of the sample can increase (particles move faster, more kinetic energy)  The state of the sample can change to a farther apart one (solid to liquid or liquid to gas, more potential energy)  Bonds can be broken (more potential energy)  Electrons move to farther energy levels o Exothermic  When energy is released, it is called an exoergic process  If it releases heat energy, it is exothermic  When energy is released, one of three things  The temperature of the sample can decrease (particles move slower, less kinetic energy)  The state of the sample can change to a closer one (gas to liquid or liquid to solid, less potential energy)  Bonds can be made (less potential energy)  Electrons move to closer energy levels Specific Heat o Specific heat- related to how much energy involved in temperature changes of matter

The amount of heat energy required to raise the temperature of 1g of a substance by 1 C without state change  S = heat absorbed per mass per C rise  S = q/m x deltaT  q = heat absorbed by a sample  if heat is released, q is negative  q = m x S x deltaT  deltaT = rise in temperature  If temperature drops, deltaT is negative o Final temperature – initial temperature  deltaT = Tfinal – Tinitial o When to use  Specific heat is only used for heat energy transfers resulting in temperature change of matter  Warming a sample, kinetic energy increases  Cooling a sample, kinetic energy decreases o Hints about problems  Only for a sample that remains in one stage (gas, liquid, solid) with no chemical change  q = mass x specific heat x temperature change  q = m x S x delta  deltaT always Tfinal – Tinitial  Look up specific heat values on a reference table  Cooling: - deltaT = -q (- means losing heat or output of heat)  Temperature change can only go up to the substances’ melting point or boiling point, then there is a state change (T will not change)  After the state has changed, the new state can change temp (but each state has a new specific heat) Calorimetry o When a hot sample transfers heat energy to a cold sample (no state change), the heat is conserved (if in a closed or adiabatic system)  Qlost + qgained = 0  Qlost = -qgained  Hot metal is placed in cold water  Mmetal (Smetal)(Tfinal-Tmetal,initial) = -mwater (Swater)(Tfinalwater,initial)  The final T is the same (thermal equilibrium)  Can solve for any one value if you know all the others 

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