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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?