Download File - Biochemistry

Biohemistry
NDHS
Name: _______________________
Date: ________________________
ATOMIC THEORY:
Greek Philosophers: 4 Elements –
Democritus: world made of two things –__________________ and tiny
particles called ______________ = unable to cut
•
atoms are the smallest particles and each substance had its own
type of atom
- wood atoms, air atoms, water atoms
Dalton:
1. all matter is made of tiny particles called atoms
2. atoms can’t be broken down further
3. atoms of different elements differ
4. atoms of the same element are identical
5. atoms combine to form compounds in specific ratios and can be
rearranged to make new compounds
Atomic Composition:
surrounded by
contains
Charges:
Protons =
Electrons =
Neutrons =
Mass:
Protons =
Neutrons =
Electrons =
Breaking Down the Nucleus:
Protons and Neutrons
Protons:
- the number of protons in the atoms of an element is ____________
- changing the element will change the ________________
The Number of Protons = The ___________________
The Number of Protons and Neutrons = The _________________
Element Symbols
Symbol, Atomic Mass, Atomic Number
- information: element, # of protons (and electrons by implication), # of
neutrons
Differences Among Atoms of the Same Element
IONS:
- in electrically neutral atoms, __________________
- loss or gain of electrons  _____
________ = __________________
________ = __________________
Ex: Sodium – atomic # =
11 _________, 11 __________
- lose 1 electron = ____________
Chlorine – atomic # = 17
17 ________, 17 _________
- gains 1 electron = __________________
=
Calculating Atomic Mass and Charge
EX:
An atom has an atomic number of 12, 13 neutrons and 12 electrons.
Identity:
Atomic Mass:
Charge:
An atom has an atomic mass of 35, 18 neutrons and 18 electrons.
Identity:
Protons:
Charge:
An atom has 20 protons, 21 neutrons and 18 electrons.
Identity:
Atomic #:
Atomic Mass:
Charge:
ISOTOPES:
- same number of protons, different numbers of _________
- changes the __________________
ATOMIC MASS= # of protons and neutrons
 each has about the same mass which is designated as an
__________________
 Determination of amu

one element chosen as a standard – Carbon 12

6 protons, 6 neutrons = 12 amu

therefore 1 amu = 1/12th Carbon atom
HOWEVER:
Is the Atomic Mass of an Element a whole number? NOPE
WHY? The Atomic Mass on the Periodic Table is the __________________
_____________________________________________________________
_____________________________________________________________
Calculating Average Atomic Mass:
[(# of atoms X mass of isotope A) + (# of atoms X mass of isotope B) + . . .]
divided by (total number of atoms of all isotopes combined)
OR
(Mass of Isotope A X Relative Abundance) + (Mass of Isotope B X
Relative Abundance) + . . . = Average Atomic Mass
Examples:
Boron has two naturally occurring isotopes: Boron-10 (abundance =
19.8%, mass = 10.013 amu) and Boron-11 (abundance = 80.2%, mass =
11.009 amu. Calculate the atomic mass of boron.
Calculate the atomic mass of magnesium. The three magnesium
isotopes have atomic masses and relative abundances as follows:
23.985 amu (78.99%)
24.986 amu (10.00%)
25.982 amu (11.01%)
Radioactive Decay
________________________:
___________________ - Radioisotopes
- atoms are unstable because they are __________________
- atoms give off the energy (______________) to become
more ______________
- Process of losing the energy is__________________
o Atoms can actually become __________________
Types of Radiation:
– positively charged _______________
Alpha particle decay: Unstable nucleus loses an alpha
particle – result, atom loses _________________________________
- mass decreases by _______
- atom becomes another element
– negatively charged particle (___________)
Beta particle decay: results from the break down of a
neutron into a __________________
- atom becomes another element
– high energy (_________________ )
Gamma Ray Emission: following Beta particle decay the
nucleus still has ___________ so the nucleus releases it as ______________
- both the atomic mass and number remain the ______
– release of a ____________________________
- the ______
but does________________ – the mass of a
______________________
____________________
_________________ – rare instance where an______________
_______________ – the _________________________________________
– therefore the _________________________________________________
Calculating Nuclear Decay:
_____________________________________________________________
Ex: Uranium-238 takes _______________
If you had a 10.0g sample, it would take 4.47x109 years for
it to decay to ________
It would take 4.47x109 more years to degrade to _____
It would take 4.47x109 more years to degrade to _____…….
Half Life 0
1
2
3
4
Time
0
10
10 more
20 total
10 more
30 total
10 more
40 total
Percent
Remaining
100%
50%
25%
12.5%
6.25%
Fraction
Remaining
1
½
¼
1/8
1/16
Equations:
Amount Remaining = (Initial amount)(1/2)n
n = number of half lives that have passed
Amount Remaining = (Initial amount)(1/2)t/T
t = elapsed time
T = duration of half-life
Ex: Radioactive iodine-131 has a half-life of 8.04 days
1. If you have 8.2 ug (micrograms) of this isotope, what mass remains
after 32.2 days?
2. How long will it take for a sample of iodine-131 to decay to 1/8 of its
activity?