2. What is "Stuff" Made of?
• Anything that occupies space and has
mass is matter.
– Matter is all the “stuff” in the physical
universe
– All matter is made of atoms joined
together.
• Matter can assume different forms:
– B.E. Condensate
– Solid: definite shape and volume
– Liquid: changeable shape, definite volume
– Gas: changeable shape and volume
– Plasma
• Mass is a measurement of how much matter an object contains
– weight
• measures gravity's pull on the atoms in the matter of an object
– density
• measures how tightly packed together the atoms of an object are
– volume
• is a measurement of how much space the atoms of an object takes up
3. Elemental Chemistry
• An element is the simplest form of matter
– each element has unique chemical properties
• how they interact with other atoms
– some elements share physical properties
• measurable or detectable with our senses
• (what it looks or smells like, its dense, etc.)
– comprised of a single type of atom
4. The Atom
• Atoms are the building blocks of matter
– 100 trillion atoms can fit on the head of a pin
• Atoms join together (bond) to form
elements
• The different parts of an atom
are called subatomic particles:
– electrons
• negative (-) charge
– protons
• positive (+) charge
– neutrons
• no charge (neutral)
5. Atomic Structure
electron shell
aka orbital
proton aka energy level
nucleus
N
+ -
N
+
-
electron neutron
HELIUM ATOM
6. Elements & Compounds
• Each element is made up of only ONE kind of atom.
– elements defined by # of protons (atomic number)
– named for their properties or after scientists (or their countries)
– sorted into a chart (listed by atomic # and grouped by energy levels)
• When different atoms join together, a compound is formed.
– Chlorine gas vs. Sodium chloride
– Hydrogen gas and pure Oxygen vs. Water
7. Periodic Table Entries
2 Atomic number
# of protons in the nucleus
80
He
4.003 Atomic mass
combined number of protons
Hg
200.59
and neutrons in the nucleus
(in atomic mass units)
number of electrons = number of protons
8. Isotopes
• An atom with more or fewer neutrons than protons is an
isotope
– doesn't change the electrical charge
– does influence the atomic mass
• Isotopes retain the same basic properties as the parent
element
9. Electron Shells (a.k.a. orbitals)
Electrons are arranged in energy levels
valence
or shells around the nucleus of an atom. electrons
18 34
Ar
39.948
Se
78.96
• the inner shell holds a maximum of 2 electrons valence
electrons
• second shell holds a maximum of 8 electrons
• third shell holds a maximum of 18 electrons
• e-'s in outer-most shell called valence electrons
10. Electron Cloud Theory
• Unlike the planetary model, electrons do not stay in a neat,
circular orbit around the nucleus.
• Electrons actually spin so fast
around the nucleus in different
areas that it's impossible to
pinpoint their exact location.
– like spinning fan blades look blurry
– darker blue areas are locations where
electrons are most likely to be found at
any given moment
11. Second
Octet Rule = atoms tend to gain, lose or share electrons so as to
have filled energy levels (usually 8 electrons)
C would like to Gain 4 electrons
N would like to Gain 3 electrons
O would like to Gain 2 electrons
12. Week 3 Lab
Water & Diffusion
Egg Permeability Experiment
13. Ions
• Atoms lose or gain an electron to become electrically charged
– loss of electron results in more protons than electrons (positive charge)
– gained electron causes more electrons than protons (negative charge)
• Usually, ions
retain the
same basic
properties as
the base
element
• Isoelectronic =
having the
same electron
configuration
14. Chemical Bonding
• Elements are attracted and “stick” to each other as they attempt to fill
electron shells
• A chemical reaction is when atoms interact, forming or breaking bonds
– Two kinds of bonds we often see in Biology are:
1. Covalent bonds – electrons are shared
2. Ionic bonds – electrons are transferred
15. Ionic Bonds
• Ionic bonds form between atoms with opposite charges
– bond formed by the transfer of electrons
– produce charged ions
– Examples: NaCl, CaCl2
Calcium
chloride
Sodium
chloride
16. Sodium (metal) + Chlorine (gas) = Table Salt?
• One electron from Sodium (Na) is transferred to Chlorine (Cl).
• This causes a charge imbalance in each atom as ions are formed.
• The Na becomes (Na+) and the Cl becomes (Cl-).
• Opposite charges of the ions cause an attraction that keeps them bonded together
• So through ionic bonding, two elements (a HIGHLY reactive metal and a DEADLY
poisonous gas) form a fairly stable compound that we eat all the time!
18. Covalent Bonds
• Covalent bonds form between elements of similar charges
– formed by sharing electron pairs
– Examples: CO2, O2, H2O Oxygen
gas
Lewis Dot Structures
Water
Carbon
dioxide
19. Non-polar Covalent Bond
With non-polar H2
molecules, two
atoms share
one or more F2
pairs of outer-
shell electrons
equally.
CO2
CF4
- small (subscript) number refers to # of atoms of PREVIOUS element in each molecule
20. Polar Covalent Bond
In polar molecules, two atoms share one or more pairs of outer-
shell (valence) electrons unequally.
- water is a polar molecule because oxygen is more electronegative than
hydrogen, and therefore electrons are pulled closer to oxygen.
21. Electronegativity of Bonds
• How badly an atom "wants" (attracts) an electron will
determine the type of bond and resulting electronegativity
(charge) of the resulting molecule.
• Ionic bonds form when atoms so
strongly want to gain and donate
electrons, an actual transfer
occurs.
• Non-polar covalent bonds form
when electrons are shared equally,
so both atoms "want" those e-'s to
the same degree.
• Polar covalent bonds form when
one atom wants an electron more
strongly than the other, but is still
willing to share, so the shared
electrons spend more time around
one atom than the other.
Notes de l'éditeur
goo·gol: a number that is equal to 1 followed by 100 zeros B.E. Condensate: super-cooled matter (atoms no longer act as individual particles) Solid: retains shape, cannot be compressed (atoms vibrate while locked in place) Liquid: assumes shape of container (atoms touch but slide past each other) Gas: fills container, compressible (atoms far apart, move freely) Plasma: super-heated gas (speedy free electrons and nuclei)
"pure" = a mass of only one element many elements are unstable and not found to be naturally occurring in nature
How Small is an Atom? A textbook or encyclopedia will tell you that a typical atom is about 0.0000001 of a millimeter in diameter. Now, how are we supposed to picture that? That information may be helpful for calculating how big a carton you'll need for mailing a given number of atoms to someone, but then again, maybe not. It certainly doesn't help the typical person understand what all the fuss is about. Here's a way you can visuallize the size of an atom, and the vast number of atoms that comprise the everyday objects in your world. For this experiment, I'd like you to get a pin. Find an ordinary sewing pin, now, before reading any further. Got one? Good. Look at the head of the pin closely. If you have one, use a magnifying glass or even a microscope to look long and closely at the head of the pin. Just stare at it for a while. Now close your eyes and imagine yourself shrinking down, almost vanishing, descending down onto on the head of the pin. You have shrunk down so small that the head of the pin is a vast desert on which you are standing, the edges of which you cannot see. You begin walking in one direction. You continue walking on the head of the pin for many days before coming within sight of the edge. How far have you walked? 50 miles? 100 miles? For the first time, you look down at the surface you have been walking on. It's fairly smooth, with an occasional ditch to stride over and variations up and down. You recognize this as the results of polishing when the pin was manufactured. You kneel down for a closer look at this shimmering surface and notice that it seems to be made up of small marbles packed closely together, and they are all vibrating slightly. Now you are seeing individual atoms. There are mostly iron, but also nickel, copper, and several other species, distinguishable by their differing sizes. Open your eyes and look again at the head of the pin. You will now be able to visuallize how small atoms are and how impossibly many there are just on the head of that pin. Now I'd like you to picture yourself on the beach. Next time you go to the beach, remember this and go through these steps. Stand on the beach and look up and down the shoreline. Picture in your mind how deep the sand might be. Deeper than a house is tall? How much sand? Reach down and pick up a handful of sand. How many grains of sand do you see? Could you even begin to count the number of grains of sand you are holding in that one handful? Allow most of the sand to fall through your fingers. Inevitably, a few grains remain clinging to your skin. Look closely at your hands, and try to pick out one single grain of sand. While looking at that one grain of sand, say the following words: "There are possibly more atoms within that single grain of sand than there are grains of sand on this entire beach." It is no wonder then that the existence of atoms was completely unknown for so long, and then, debated for so long. By about the start of the 20th century, the indirect evidence of atoms from observations made over the previous few hundred years had finally won over most people to the concept and existence of atoms. Now, we have technology that can directly image the atoms on the surface of a grain of sand or on the head of a pin.
Simplest form of matter (for unique chemical properties) = an element
Planetary model vs. Electron Cloud theory Inner levels require less energy for orbit, outer considerered "higher energy state"
