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BIOL 1F90 Midterm Study Guide- Chapter 2 .docx
BIOL 1F90 Midterm Study Guide- Chapter 2 .docx

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Brock University
Douglas Bruce

BIOL 1F90 Midterm: Study Guide Chapter 2: The Chemical Basis of Life I: Atoms, Molecules and Water Atoms  Smallest functional units of matter that form all chemical substances  Can’t be further broken down into other substances by ordinary chemical of physical means.  Chemical element: A specific type of atom Table 2.1: Characteristics of Major subatomic Particles Particle Location Charge Mass relative to electron Electron Around the -1 1 nucleus Proton In the nucleus +1 1,836 Neutron In the Nucleus 0 1,839 The number of protons and electrons are identical, but the number of neutrons vary. Atomic Nucleus: Centre of the atom where protons and neutrons are confined.  Entire atom has no net electric charge because the positively charged protons tend to attract an equal number of negatively charged atoms Rutherford Determined the Planetary Model of the Atom  Hypothesis: Atoms in gold foil are composed of diffuse, evenly distributed charges that should usually cause alpha particles to be slightly deflected as they pass through  Fine beam of positively charged alpha particles (helium nuclei) shoots through a thin sheet of gold o If the positive charges of the gold atoms were uniformly distributed, most of the alpha particles would be slightly deflected o The most important feature of the electric charge is like charges repel each other.  Conclusion: Most of the volume of an atom is empty space , with the positive charges concentrated in a small volume Data Collected % of alpha particles detected on Location zinc sulphide screens 98% Undeflected <2% Slightly deflected 0.01% Bounced back  98% passed through; concluded that most of the volume of an atom is empty space  0.01 bounced back; assumes that most of the atom’s positive charge is localized in a highly compact area  99.99% of an atom’s volume is outside the nucleus. Orbitals:  A place electrons are likely to be found orbiting around the nucleus  Cloudlike orbitals; the cloud represents the region where the electron would most likely to be found.  s orbitals: Orbitals that are spherical  p orbitals: Orbitals that are shaped like dumbbells or propellers  A orbital can contain a maximum of two electrons; any atom with more than 2 electrons must contain more than one orbital.  Energy shells are in the orbital, associated with specific energy.  The first energy shell is closest to the nucleus  Atoms with progressively more electrons have orbitals within electron shells  Are at greater distances from the center of the nucleus o First shell is 1 spherical orbital (1s) which holds 2 electrons o 2 ndshells is 1 spherical orbital (2s) and 3 dumbbell shaped orbitals ( 2p)  Valance electrons: Electrons in the outer shell o Available to combine with other atoms o Involved in sharing and carrying electrons  Electron Energy: An atoms electrons vary in the amount of energy they possess. Energy is defined as the capacity to cause change.  Potential Energy: the energy that matter possesses because of its location or structure Example: Nitrogen  The first energy shell is filled with 2 electrons while the second energy shell with 5 electron  Since the second energy shell can fill up to 8 electrons, this shell is now fully filled yet.  Atoms with unfilled energy shells tend to share, release, or obtain electrons to fill their outer energy shell. Protons  Atomic Number: The number of protons in an atom o Ie. The atomic number for hydrogen is 1  Hydrogen have 1 proton  Atoms are electrically neutral. Therefore, the atomic number also indicates the number of electrons in the atom.  Atomic number and electron energy is useful for organizing chemical elements into the periodic table Periodic Table  Each row, top to bottom corresponds to the number of electron shells  Columns, left to right, indicate the number of electrons in the outer shell  Reasons for similarities of elements within a column o Same number of electrons in their outer shell o Similar chemical bonding properties Atomic Mass  Measured in units called Daltons (Da),named after the chemist John Dalton who believed that matter is composed of minute, indivisible units he called atoms  One Dalton ( Da)= ½ the mass of a carbon atom  Mole: A conversion of grams  daltons o Contains the same number of atoms as are present in exactly 12 g of carbon o 12g of carbon = 1 mole  1 mole of any element contains o The same number of atoms as in exactly 12g of carbon  The most common form of carbon has 6 protons and 6 neutrons  exactly 12 daltons o Hydrogen atom  1 dalton, 1/12 the mass of a carbon atom o Magnesium atom  24 daltons, twice the mass of a carbon atom 23 o Avogadro’s Constant: 6.022 x 10 atoms Isotopes  Many elements can exist in multiple forms  Differs in the number of neutrons they contain  Ie.12C, contains 6 protons and 6 neutrons = atomic mass of 12 daltons o 99% of living organisms contains 12C o 14C, a rare carbon contains 6 protons and eight neutrons = atomic mass of 14 Daltons  Thus, the average atomic mass of carbon is very close to, but slightly greater than, 12 Daltons because of the existence of a small number of heavier isotopes  Radioisotopes: Isotopes that are unstable and found in nature o They lose energy through radiation by emitting subatomic particles o Modern medicine makes use of the high energy level of radioisotopes in many ways. Eg. Solutions containing radioisotopes of iodine can be given to a person with an overactive thyroid gland o PET scans Carbon, Hydrogen, Oxygen, Nitrogen  Account for the vast majority of atoms in living organisms  These elements typically make up about 95% of the mass of living organisms  Hydrogen and oxygen are primarily in water o 60% of the mass of most animals o 95% or more in some plants  Nitrogen and carbon ( along with hydrogen and oxygen) are major building blocks of all living matter.  Mineral elements: less than 1% ie. Calcium and phosphorus, are important constituents of the skeletons and shells of animals  Sodium and potassium are key regulators of water movement and electrical potentials that occur across membranes.  Trace elements: present in extremely small quantities, but still essential for normal growth and function Chemical Bonds and Molecules  The linkage of atoms with other atoms serves as the basis for life and gives life its great diversity  Molecule: 2 or more atoms bonded together  Molecular formula: Consists of chemical symbols for all the atoms present and subscripts that show the number of atoms that are present in the molecule  Compound: A molecule made of 2 or more different elements o Most important features of compounds is the new properties that emerge, meaning that the properties of a compound can differ greatly from the elements that combined to form it. Covalent Bonds  Atoms share a pair of electrons  Occurs between atoms whose outer shells are not full.  They are strong chemical bonds because the shared electrons behave as if they belong to both atoms.  In some molecules, a double bond occurs when atoms share 2 pairs of electrons rather than one pair  Triple Bonds may form, when 3 pairs of electrons are shared between 2 atoms  Octet Rule:  atoms are stable when their outer shell is full  for many atoms, the outer shell fills with 8 electrons  one exception is hydrogen which fills its outer shell with 2 electrons Remembering the number of unshared electrons ( number of covalent bonds ) for these four atoms really makes understanding and learning the structure of biological important molecules easier* Electronegativity  A measure of its ability to attract electrons in a bond from another atom  When two atoms with different electronegativity form a covalent bond, the shared electron’s orbit is most likely to be closer to the atom of higher electronegativity than to the atom of lower electronegativity, creating a polarity in the distribution of the charge  The more electron negative an atom, the more strongly it pills shared electrons toward itself  Ie. Oxygen: Oxygen is a pig for electrons and nitrogen is almost as bad. They don’t share equally with carbon or hydrogen Polar Covalent Bonds ( Dipole)  Skewed distribution of electrons toward one of the two atoms creates a polarity across the molecule  Ie. 2 O  a classic example of a polar covalent bond o The shared electrons tend to be closer to oxygen than hydrogen atoms because oxygen has a higher electronegativity o A partial negative charge (oxygen) and 2 partial regions of partial positive charge (hydrogen) Non Polar Covalent Bonds  Bonds between atoms with similar electronegativities  Equal sharing of electrons  Molecules formed are called non polar molecules  Molecules can have same regions with non polar bonds and other regions with polar bonds Ionic Bonds  Ion: When an atom or a molecule gains or loses one or more electrons (with no change in number of protons), it acquires a net electric charge  Cations: Ions that have a net positive charge  Anions: Ions that have a net negative charge  Ionization: The process of ion forma
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