Chapter 2 Matter and Energy

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Chapter 2Matter and Energy
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Representations of Matter

• Goal 1
• Identify and explain the differences among
  observations of matter at the macroscopic,
  microscopic, and particulate levels.
• Goal 2
• Define the term model as it is used in chemistry
  to represent pieces of matter too small to see.

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Representations of Matter

• Anything that has mass (sometimes expressed as
  weight) and takes up space is called matter.
• Matter can be observed and/or thought about at
  different levels
• Macroscopic
• Microscopic
• Particulate

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Representations of Matter

• Macroscopic Samples of Matter
• Mountains
• Rocky cliffs
• Huge boulders
• Rocks and stones
• Gravel
• Sand
• Macro- means large

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Representations of Matter

• Microscopic Samples of Matter
• Tiny animals or plants
• Cells
• Crystals on rock surfaces
• Micro- means small

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Representations of Matter

• Particulate Samples of Matter
• Too small to see, even with the most
• powerful optical microscope.
• Chemists imagine the nature of the behavior of
  the tiny particles that make up matter, based on
  experimental measurements and models of matter,
  and they use that knowledge to carry out changes
  from one type of matter to another.

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Representations of Matter

• Macroscopic, Microscopic, and Particulate Matter

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Representations of Matter

• Model
• A representation of something.
• Geologists model the earth (globe).
• Biologists model cells.
• Chemists model atoms and molecules.

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Representations of Matter

• Ball-and-Stick Model
• Symbolizes atoms as balls and the electrons that
• connect those atoms as sticks.
• Space-Filling Model
• Shows the outer boundaries of the particle
• in three-dimensional space.

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Representations of Matter
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Representations of Matter

• Models are represented in writing with symbols.
• Chemical symbols are letters that represent atoms
  of elements.
• H represents an atom of hydrogen.
• O Represents an atom of oxygen.
• H2O represents a molecule of water
• Two hydrogen atoms and one oxygen atom.

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Representations of Matter

• Chemists make mental transformations between
  visible macroscopic matter (the liquid water) and
  models of the particulate-level molecules that
  make up the matter (red and white space-filling
  models). Written symbols (HOH on the board)
  serve as simpler representations of the
  particulate-level models.

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States of Matter

• Goal 3
• Identify and explain the differences among gases,
  liquids, and solids in terms of (a) visible
  properties, (b) distance between particles, and
  (c) particle movement.

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States of Matter

• States of Matter
• Familiar examples of the states of matter
• The air you breathe is a gas.
• The water you drink is a liquid.
• The food you eat is a solid.

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States of Matter

• Kinetic Molecular Theory
• All matter consists of extremely tiny particles
• that are in constant motion.
• Kinetic refers to motion.
• Molecular comes from molecule, the smallest
• unit particle that can exist independently and
• possess the identity of the substance.
• Theory is a hypothesis that has been tested and
• confirmed by many experiments.

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States of Matter

• The speed at which particles move is faster at
  higher temperatures and slower at lower
  temperatures.
• There is an attraction among particles in all
  samples of matter.
• The state of matter of any sample depends on
  temperature (the speed at which particles move)
  and the attractions among the particles that make
  up the sample.

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States of Matter

• Particulate-Level Behavior of States of Matter
• Gas
• Particles are independent of one another,
• moving in random fashion
• Liquid
• Particles move freely among themselves, but clump
  together
• Solid
• Particles vibrate in fixed positions relative to
  one another

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States of Matter
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Properties and Changes

• Goal 4
• Distinguish between physical and chemical
  properties at both the particulate level and the
  macroscopic level.
• Goal 5
• Distinguish between physical and chemical changes
  at both the particulate level and the macroscopic
  level.

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Properties and Changes

• Physical Properties
• Description as detected by senses
• Color, shape, odor, etc.
• Measurable properties
• Density, boiling point, etc.
• Examples
• Charcoal is black
• Glass is hard
• The normal boiling point of water is 100C

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Properties and Changes

• Physical Changes
• Alteration of the physical form of matter
• without changing its chemical identity
• No new substance formed
• Examples
• Ice melts to liquid water
• Dry ice changes to gaseous carbon dioxide
• A rock is ground into sand

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Properties and Changes

• In a Physical Change, the Molecules are Unchanged

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Properties and Changes

• Chemical Changes
• Chemical identity of a substance is destroyed
• A new substance is formed
• Examples
• Water decomposes to hydrogen and oxygen gases
  when subjected to an electrical current
• Iron rusts
• Food is digested

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Properties and Changes

• In a Chemical Change, the Molecules Change

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Properties and Changes
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Properties and Changes

• Chemical Properties
• The types of chemical change a substance
• is able to participate in.
• Examples
• A chemical property of water is that it can be
  decomposed to its elements when subjected to an
  electrical current.
• A chemical property of iron is that it will
• rust under certain conditions.
• A chemical property of starch is that it reacts
  to form
• sugar during digestion.

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Properties and Changes
Chemical Physical
Changes Chemical identity of a substance is destroyed New substances formed New form of same substance No new substances formed
Properties Types of chemical changes possible Description as detected by the senses Measurable properties
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Substances and Mixtures

• Goal 6
• Distinguish between a pure substance and a
  mixture at both the macroscopic level and the
  particulate level.
• Goal 7
• Distinguish between homogeneous and heterogeneous
  matter.

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Substances and Mixtures

• Pure Substance
• A sample consisting of only one kind of matter,
• either compound or element
• made up entirely of one kind of particle.
• Unique set of physical and chemical properties.
• Cannot be separated into parts by a physical
  change.

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Substances and Mixtures

• Mixture
• A sample of matter that consists of two or more
  substances.
• Physical and chemical properties of a mixture
  vary with
• different relative amounts of the parts.
• Can be separated into parts via physical
  processes.

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Substances and Mixtures

• Pure water has a constant boiling pointa
  physical property.
• The boiling point of a mixture (solution) changes
  as the composition of the mixture changes.

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Substances and Mixtures

• You cannot distinguish a pure substance from a
  mixture of uniform appearance by observation
  alone at the macroscopic level.

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Substances and Mixtures

• Solution
• A homogeneous mixture of two or more components.
• Homogeneous
• A sample that has uniform
• appearance and composition throughout.
• Examples Tea, paint, gasoline
• Heterogeneous
• A sample with different phases, usually visible.
• Examples Carbonated beverages, salad dressings

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Substances and Mixtures

• Homogeneous Matter may be Either a Pure Substance
  or a Mixture

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Separation of Mixtures

• Goal 8
• Describe how distillation and filtration rely on
  physical changes and properties to separate
  components of mixtures.

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Separation of Mixtures

• Most natural substances are mixtures.
• Separation processes are an important part of
  chemistry.
• Nitrogen and oxygen are separated from the
  mixture called air.
• Pure water is separated from the mixture called
  natural water.
• Gasoline is separated from the mixture called
  crude oil.

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Separation of Mixtures

• A Physical Property, Magnetism, Allows a Mixture
  of Iron and Sulfur to be Separated

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Separation of Mixtures

• Distillation
• Separation of the parts of a mixture by heating a
  liquid solution until one component boils,
  changing into the gaseous state.
• The pure substance in the gaseous state is then
• collected and cooled into the liquid state.
• Boiling is a physical change.
• Distillation allows components in a homogeneous
  mixture
• to be separated into one or more pure substances.

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Separation of Mixtures

• Laboratory Distillation Apparatus

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Separation of Mixtures

• Filtration
• Separation of the components of a mixture by
  physical means by using a porous medium, such as
  filter paper, to separate components based upon
  relative particles sizes.
• Filtration is based on the physical properties of
  a mixture
• The particle sizes of a component to be separated
• must be significantly larger or smaller than the
• pore size of the filtration medium.

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Separation of Mixtures

• Gravity Filtration

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Elements and Compounds

• Goal 9
• Distinguish between elements and compounds.
• Goal 10
• Distinguish between elemental symbols and the
  formulas of chemical compounds.
• Goal 11
• Distinguish between atoms and molecules.

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Elements and Compounds

• Element
• Pure substance that cannot be decomposed into
  other pure substances by ordinary chemical means.
• Atom
• Smallest particle of an element that can combine
  with atoms of other elements to form chemical
  compounds.
• Compound
• Pure substance that can be broken down into two
  or more
• other pure substances by a chemical change.

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Elements and Compounds

• The Element Silver and a Particulate-Level Model
  of Silver Atoms

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Elements and Compounds

• Mixtures are separated into pure substances by
  physical means compounds are separated into pure
  substances by chemical changes.

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Elements and Compounds

• Elements
• At least 88 elements occur in nature.
• Examples copper, sulfur, gold, silver
• 11 elements occur in nature as gases
• 2 occur as liquids (mercury and bromine)
• the others occur as solids.
• Name of an element is always a single word
• compound names are usually two words
• or a polysyllabic compound word.

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Elements and Compounds

• Major Elements of the Human Body
• Element Percentage Composition by Number of Atoms
• Hydrogen 63.0
• Oxygen 25.5
• Carbon 9.45
• Nitrogen 1.35
• These four elements make up 99.3 of the atoms in
  your body.

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Elements and Compounds

• Familiar Objects that are Composed of Nearly Pure
  Elements

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Elements and Compounds

• Familiar Objects that are Compounds

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Elements and Compounds

• Elemental Symbols
• Letters that symbolize elements
• The first letter of the name of the element,
• written in uppercase, is often its symbol.
• Examples Hydrogen, H
• Oxygen, O
• Carbon, C
• If more than one element begins with the same
  letter,
• a second lowercase letter is added.
• Examples Helium, He
• Osmium, Os
• Chlorine, Cl

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Elements and Compounds

• Chemical Formulas
• Symbolic representations of the particles of a
  pure substance.
• A combination of the symbols of all the elements
  in a substance.
• The formula of most elements is the same as the
  symbol of the element, e.g., helium He sodium,
  Na.
• Other elements exists in nature as molecules and
  their formulas indicate the number of atoms of
  the element in the molecule, e.g., hydrogen, H2
  oxygen, O2.

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Elements and Compounds

• Formula Unit
• Molecule or simplest ratio of particles for
  non-molecular species.
• Ammonia molecules have the formula NH3
• 1 atom of nitrogen and 3 atoms of hydrogen.
• Magnesium chloride exists as an orderly,
  repeating pattern
• of magnesium and chlorine in a 12 ratio
• Its formula unit is MgCl2.

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Elements and Compounds

• Law of Definite Composition or
• Law of Constant Composition
• Any compound is always made up of elements in
• the same proportion by mass (weight).
• No matter its source, water (H2O) is
• 11.1 parts hydrogen per 88.9 parts oxygen.

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Elements and Compounds

• The Properties of a Compound are Different from
  the
• Properties of the Elements that Make Up the
  Compound
• Water, H2O
• Liquid at 25C, melts at 0C, boils at 100C
• Hydrogen, H2
• Gas at 25C, melts at 259C, boils at 253C
• Oxygen, O2
• Gas at 25C, melts at 219C, boils at 183C

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Elements and Compounds

• Particulate and Macroscopic Views of Elements and
  Compounds

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Elements and Compounds

• Particulate and Macroscopic Views of Elements and
  Compounds

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Elements and Compounds

• Particulate and Macroscopic Views of Elements and
  Compounds

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Elements and Compounds

• Summary of the Classification System for Matter

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Electrical Character of Matter

• Goal 12
• Match electrostatic forces of attraction and
  repulsion with combinations of positive and
  negative charge.

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Electrical Character of Matter

• Two of the fundamental forces that govern
• the operation of the universe are
• Force of gravity
• Electromagnetic force
• The electromagnetic force plays an important role
  in understanding chemistry.
• It includes electricity and magnetism.

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Electrical Character of Matter

• Force Field
• Region in space where the force is effective.
• Electrostatic Force
• The force of an electrical charge that does not
  move.
• A charged object exerts an invisible
  electrostatic force.

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Electrical Character of Matter

• Experimental evidence shows that
• There are only two types of electrical charge,
• positive and negative.
• Two objects having the same charge,
• both positive or both negative, repel each other.
• Two objects having unlike charges,
• one positive and one negative, attract each other.

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Electrical Character of Matter

• Electrostatic forces show that matter has
  electrical properties.
• These forces are responsible for the energy
  absorbed or released in chemical changes.

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Chemical Change

• Goal 13
• Distinguish between reactants and products in a
  chemical equation.
• Goal 14
• Distinguish between exothermic and endothermic
  changes.
• Goal 15
• Distinguish between potential energy and kinetic
  energy.

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Chemical Change

• Chemical Equation
• A symbolic representation of chemical change,
  with the formulas of the beginning substances to
  the left of an arrow that points to the formulas
  of the substances formed.
• Reactant
• Original substance
• Product
• Substance formed as a result of chemical change
• 2 H2O 2 H2 O2
• Reactant Products

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Chemical Change

• Exothermic Reaction
• A chemical change that releases energy to its
  surroundings.
• ExampleBurning charcoal
• C O2 CO2 energy

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Chemical Change

• Endothermic Reaction
• A chemical change that absorbs energy from its
  surroundings.
• Example
• Decomposition of water to its elements
• 2 H2O energy 2 H2 O2

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Chemical Change

• Energy
• The ability to do work.
• Potential Energy
• Energy due to position in a field where forces of
• attraction and/or repulsion are present.
• Gravitational Potential Energy
• Position in the earths gravitational field.
• Electrical Potential Energy
• Position in an electrical field.

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Chemical Change

• Minimization of energy
• One of the driving forces that cause chemical
  reactions to occur.
• Chemical energy comes largely from the
  rearrangement of
• charged particles in an electrostatic field.

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Chemical Change

• Kinetic Energy
• Energy of motion
• The temperature of an object is proportional to
  the
• average kinetic energy of its particles.

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Conservation Laws

• Goal 16
• State the meaning of, or draw conclusions based
  on, the Law of Conservation of Mass.
• Goal 17
• State the meaning of, or draw conclusions based
  on, the Law of Conservation of Energy.

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Conservation Laws

• The Conservation Law
• In any change, the sum of mass plus energy is
  conserved
• they are neither created nor destroyed.
• ?E ?m ? c2
• Matter is an extremely concentrated form of
  energy.

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Conservation Laws

• Law of Conservation of Mass
• In a nonnuclear change, mass is conserved
• it is neither created nor destroyed.
• Mass may change form, however.
• In any ordinary chemical change,
• Total mass of reactants Total mass of products

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Conservation Laws

• Law of Conservation of Energy
• In a nonnuclear change, energy is conserved
• it is neither created nor destroyed.
• Energy may change form, however.
• The energy lost in one form is always exactly
  equal
• to the energy gained in another form.

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Conservation Laws

• Common Events in which Energy Changes
• from One Form to Another