the nature of matter
DESCRIPTION
The Nature of Matter. Chapter 1 BIOLOGY 391. What is everything made of?. MATTER Anything that has mass and takes up space ATOM The smallest unit of matter. ATOMS. Basic unit of matter Size: 1,000,000 (million) side by side = 1 cm Atoms like to be neutral- no charge - PowerPoint PPT PresentationTRANSCRIPT
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The Nature of Matter
Chapter 1BIOLOGY 391
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What is everything made of?
• MATTER Anything that has mass and takes up space
• ATOM The smallest unit of matter
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ATOMS• Basic unit of matter
• Size: 1,000,000 (million) side by side = 1 cm
• Atoms like to be neutral- no charge
– Equal number of protons and electrons
– There are specific numbers of “sub-atomic” particles that the atom wants• Special cases are called isotopes or ions
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STRUCTURE OF ATOM
– NUCLEUS - protons and neutrons held together by the “strong force”
• Protons (+)• Neutrons (o)
– ELECTRON CLOUD (orbitals) – electrons surround nucleus • Electrons (-) only contains about 1/1836 the
mass of proton or neutron• Constantly moving within orbital- attracted to
the nucleus by the “weak force”
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ATOMS •The smallest particle of an element that can exist and still have the properties of that element.•An atom has many subatomic particles
Particle Charge Location Mass
Proton + Nucleus 1 amu *
Neutron N Nucleus 1 amu
Electron - ElectronCloud
0 **
•Atomic mass unit or Dalton
•** The mass of an e- is so small it’s negligible.
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Electron orbitals
1st orbital can only hold 2 electrons (too close to nucleus- not much space)
2nd orbital can hold up to 8
3rd orbital can hold up to 8
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How many atoms are there?PERIODIC TABLE
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Atoms are grouped as Elements
• Are listed on the Periodic Table• Dimitri Mendele’ev 1869• Arranged in order of their atomic #’s• Table is divided into Groups and Periods• The atomic number and atomic mass are
given for each element. Example: 6
C12.01
Atomic #
Atomic Mass
Element
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“Groups”Vertical Column
Numbered 1-8All elements of same group have the same # of e-’s in their valence shell & have similar chemical properties
Atomic # increases by 1 from left to right Example:
H, Li, Na, K, Rb, Cs, FrAll have 1 e- in valence (outer) shell
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“Periods”• Is a horizontal row of elements• There are 7 periods• Elements in a period have the same #
of energy levels
Example:H and He are in period 1They have 1 energy level
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The Elements are often described as Metallic and Nonmetallic
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Atomic Number# of protons
(and also # of electrons)
Chemical symbol
Name of Element
Atomic MassThe weight Of carbon
atom oraverage
weight of all isotopes
6
CCarbon
12.011
What is the difference between atoms?
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Special Form: ISOTOPE
Atoms of the same element containing different numbers of neutrons in the nucleus
– Some give off radiation – used to:• Trace atoms through a reaction or an organism• Treat cancer• Date very old, once living organisms
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2 examples of Isotopes
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Recap• What are the three subatomic particles and
their charges?• What is the only actual difference between
gold and mercury? • What is the atomic mass of lead?• What is an isotope?
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Organization of Matter
• Atoms usually do not occur alone, but exist with other atoms as:
–Elements (all of the same atoms)–Molecules or Compounds
• Same or different atoms bonded
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MATTER: anything that has mass & takes up space
ProtonsNeutronsElectrons
Does not contain C-C bonds
LIFE!
Organic
Molecules orCompounds
Inorganic
C-C, C-H bond
ATOMSElements: Shown in
Periodic Tablepure
bonded
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Elements
• A substance which cannot be split into simpler substances by a chemical rxn.
• A grouping of the same type of atoms– ORDER MATTERS!
• More than 100 elements exist (shown in the periodic table)
• Carbon Elements:
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Elements of Life
92 naturally occurring elements
Elements Found in Living OrganismsN CHOPS (macronutrients)
C HOPKINS Ca Fe Mg B Mn Cu Cl Mo Zn
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MATTER: has mass & takes up space
ProtonsNeutronsElectrons
Does not contain C-C bonds
LIFE!
Organic
Molecules orCompounds
Inorganic
C-C, C-H bond
ATOMSElements: Shown in
Periodic Tablepure
bonded
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Why do atoms bond?• An atom wants to have a complete outer shell
of electrons.
To do this, it can…• Share electrons with another atom• Give away its extra electrons• Steal extra electrons
• Bonds store ENERGY
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Remember: an atom is when its outer orbital is filled
For the “smaller” elements, the outer shell holds 8. The electron clouds Increase in size as you go across the periodic table. HOWEVER, the trend stays the same for the number needed to complete the outermost shell.
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BONDING: depends on number of electrons in the outermost orbital
shell• Covalent Bond- atoms share a pair of electrons
sometimes share 2 (double bond) or 3 (triple bond) pairs
Ionic Bond- One atom (very unstable) gives 1, 2 or 3 electrons away to another atom. The atom that loses electrons becomes positively charged. The atom that gains the electrons becomes negatively charged. The opposite charges cause the atoms to “bond” together
(opposites attract).
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Molecules bonded atoms
“Molecule” is often used to refer to an individual grouping. “singular”
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Compounds: also, 2 or more atoms bonded together.
Often referred to as a larger conglomerate of bonded molecules
• Order matters!• These 2 cmpnds are
made of the same 3 atoms, but in a different arrangement.
• The arrangement is responsible for their affects.aspirin sucrose
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Bonding• IONIC• COVALENT
– Polar– Nonpolar
• VAN DER WAALS• HYDROGEN
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The periodic table shows the pattern of electrons each element has in its
outermost shell
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Ions • Atoms that have lost or gained electrons
cation – positively charged ion – lost electronsanion – negatively charged ion – gained electrons
• Ionic bonds are weaker than covalent bonds– Hold less energy in the bond
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Cations versus Anions• Atom that loses electrons• Positively charged ion• Elements in Groups 1, 2,
and 3 tend to lose electrons• Metallic elements tend to
form positive ionsExample:
Ca Lose 2 electrons+2 charge
• Atom that gains electrons• Negatively charged ion• Elements in groups 5, 6, and
7 tend to gain electrons• Nonmetallic elements tend
to form anionsExample:
ClGain 1 electron-1 charge
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IONIC BONDING
Na (sodium) is very unstable because it only has one e- in its outer orbital. Cl’s (chlorine) outer orbital is
almost filled. Na gives its lonely e- to Cl.
Na become Na+ Cl becomes Cl-
Their opposite charges cause them to be attracted to one another- This is an ionic bond.
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Ionic Compounds• Metals react with nonmetals forming ionic compounds• Salts• Held together by electrostatic forces
Example: + attracted to –• Most are crystalline solids at room temperature• When dissolved in water they conduct electricity
Ex. Sodium Chloride or Table SaltNa+ + Cl-
• Dissociate(break apart) in water, producing free ions
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Electrolyte• A solution that conducts electricity.• Term for salts, specifically ions.
• The term electrolyte means that this ion is electrically-charged and moves to either a negative (cathode) or positive (anode) electrode
• Ions that move to the cathode (cations) are positively charged
• Ions that move to the anode (anions) are negatively charged
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Covalent Bonds• Bonds formed when atoms share electrons• Atoms with 4 or 5 e-’s tend to share • Each e- spends part of its time around one
nucleus and then around the other nucleus.
• Sharing e-’s completes the valence shell
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Example of a Covalent BondExample of Covalent Bonding- Water
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Covalent Bonds• Bonds between non-metals
• Poor conductors of electricity
• Do not dissociate easily in water
• Two Types:Polar: Unequal SharingNon-Polar: Equal Sharing
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Van der Waals forces• Attraction between oppositely charged areas
of adjacent molecules• Remember, e- are constantly swarming
around the nucleus. At times, there may be a moment of asymmetry – All of the electrons might be on one side– This sets up a “dipole”
weaker than covalent bonds and ionic bonds
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RECAP• What is an ion?• Why is it important that atoms bond?• What causes atoms to bond?• Explain the difference between an ionic bond
and a covalent bond• What are Van der Waals forces?
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BONDS
IONIC
Transfers/takes electronsPositive/Negative chargesWeak bondEx: NaCl (salt)
COVALENT
Shares electrons NeutralStrong bondEx: H2O, CO2, NH3
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WATER – Why so Great?
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Water!• What is the chemical formula for water?• How much water covers the Earth?• How much of your body is water?• Is there water in food?• How long could you live without water?
• H2O• 75% • 60-70%• Yes!• 3 days
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Properties of Water
• Phases: Solid, Liquid, Gas• Polarity• Hydrogen bonds
– Adhesion– Cohesion
• Making Mixtures– Solutions– Suspensions
• Making Acids and Bases
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Water Density
• Ice is less dense than liquid water• When water freezes air is trapped within the
frozen ice making the cube larger and less dense
• Benefits:– Fish and plant life can survive in liquid layers of
water under ice
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PHASE CHANGES: the closeness and speed of the compounds
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Polarity
• Water is polar
• Although the compound is neutral overall there is a shift of charge within the compound
The much larger atom, Oxygen, pulls more on the shared e-
This end of the compound becomes slightlymore negative.
Hydrogen ends become slightly
positive
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Hydrogen Bonding• Due to polarity, water compounds attract to
one another• Slightly negative oxygen attracts slightly
positive hydrogen from another compound• This attraction among water is COHESION.• Water is also attracted to other materials.
This is ADHESION.
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COHESION
Water compounds attractTo glass molecules
And form a meniscus
Water compounds attract To one another-
causes water to “bead”
ADHESION
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The greatest solvent on Earth!
• Water’s polarity allows it to break ionic bonds of other compounds…creating free ions.
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Mixtures Two or more elements physically mixed together but
not chemically combined (not bonded)
1. SOLUTIONS- a solute is dissolved into a solvent – Distributes evenly– “Like dissolves Like”– Ex: Koolaid, salt water
2. SUSPENSIONS- added substance does not dissolve but breaks into small enough pieces that it remains suspended in the water and does not settle out.- Ex: blood
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RECAP
• Why does ice float on a lake?• Explain the polarity of water – how are the
charges distributed?• What is the difference between adhesion and
cohesion?• Explain the difference between a solution and
a suspension
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Water Dissociation
• Water can break apart on its own into 2 charged ions– Hydrogen ions and hydroxide ions
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Acids and Bases
Water can react to form individual ions:
H2O H+ + OH-
• In pure water this occurs naturally but the amount of H+ is always = to the amount of OH- so water remains neutral
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pH scale: “the power of Hydrogen”
• Some solutions made with water become acidic or basic. This is determined by the amount of H+ (hydrogen ions) in the solution
• pH = - log [H+]– Example: [H+] = 1 x 10 -5 pH = 5 (acid)
[H+] = 1 x 10 -9 pH = 9 (base)Remember how exponents work:
0.00001 is greater than 0.000000001
• b/c it’s logarithmic, each pH unit represents a tenfold difference in concentration of H+ ions
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Acids
• pH range from 0 – 6.99 • Any compound that forms H+ ions in solution• H+ ions > OH- ions• The closer to 0 the more acidic the solution• Examples: stomach acid, lemon juice
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Bases (Alkaline)
• pH ranges from 7.01 to 14• Any compound that forms 0H- ions in solution• OH- ions > H+ ions• The closer to 14 the more basic the solution• Examples: lye, bleach, oven cleaner
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ACID:Any compound that forms H+ ions in solution
BASE: Any cmpnd that forms 0H- ions in solution
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pH and Living Things• pH values in living cells are usually kept
between 6.5 and 7.5– Optimal pH for chemical reactions to take place in
the body– Any switch in pH could cause serious/fatal
problems
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Buffers
• Weak acids or bases that can react with strong acids or bases
• Used to regulate pH and prevent sharp sudden changes in pH
• There are natural buffers in your blood that keep the pH at 6.5 to7.5
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RECAP
• What makes a solution acidic or basic?• How is acidity measured?• A solution with pH 8.5 is considered….
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Forming…Acid Rain…Reacting with stone