Lecture 2 - The Chemical Context of Life

Lecture 2 Chapter 2: The Chemical Context of Life  Elements & Compounds Atoms o Atomic number and Mass OVERVIEW o Isotopes o Energy levels o Electron Distribution Chemical Bonding o Types...

Lecture 2 Chapter 2: The Chemical Context of Life  Elements & Compounds Atoms o Atomic number and Mass OVERVIEW o Isotopes o Energy levels o Electron Distribution Chemical Bonding o Types Chemical Reactions ELEMENTS & COMPOUNDS Matter – anything that takes up space and has mass Makes up all organisms Many forms (rocks, metals, oils, gases, living organisms) Made up of elements ELEMENTS & COMPOUNDS Elements – substances that can’t be broken down into other substances by chemical reactions 92 elements occur in nature Each element has a symbol Examples: Gold (Au), copper (Cu), carbon (C) ELEMENTS & COMPOUNDS Compound – substance consisting of +2 different elements in a fixed ratio Example: Sodium chloride (NaCl) consists of sodium (Na) and Chlorine (Cl) 1:1 ratio of metal and gas to produce table salt due to emergent properties Another example: Water (H2O) 2:1 ratio of Hydrogen (H) and ELEMENTS & COMPOUNDS Essential Elements – required by all living organisms for normal cellular functions ~20-25% of the natural elements 96% of all living matter consists of Oxygen (O), Carbon (C), Hydrogen (H), and Nitrogen (N) 3.5% of mass consists of Calcium (Ca), Phosphorous (P), Potassium (K), and Sulfur (S) ELEMENTS & COMPOUNDS Trace Elements – required by all living organisms for normal cellular functions in very small quantities Makes up ~0.5% of living organisms Examples: Iron (Fe), Iodine (I), Zinc (Zn) ATOMS Atom – smallest unit of matter Retains the properties of an element ~5 million Hydrogen atoms could fit into a pin head Symbolize atoms with the same abbreviation used for the element made up of those atoms Example: (C) stands for the element Carbon and a single carbon atom ATOMS Subatomic Particles – small pieces of matter that make up atoms Neutrons – no electrical charge (neutral) Protons – positive electrical charge (+) Electrons – negative electrical charge (-) Atomic Nucleus – dense core at the center of an atom that tightly packs together protons and neutrons Protons give the nucleus a positive charge Electrons rapidly move around the ATOMS Subatomic Particle Characteristics Neutrons and protons weigh about 1.7 x 10-24 grams (or 1 Dalton) Atomic mass unit (amu) is a Dalton Electrons weigh ~1/2,000 of a Dalton and is disregarded in calculations ATOMS Atomic Number – number of protons unique to each element Written as a subscript to the left of the element symbol Assume equal number of protons and electrons unless otherwise indicated Mass Example: 2He– total number of protons and Number neutrons in the atomic nucleus Deduce number of neutrons Written as a superscript to the left of the element symbol Example: 4He ATOMS Atomic Mass – total mass of an atom Example: 22.9898 for Na Due to isotopes ATOMS Isotopes – different atomic forms of the same element Different amount of neutrons Changes the atomic mass Most naturally-occurring elements exist as a mixture of its isotopes Example: Carbon has three naturally- occurring isotopes All three isotopes have 6 protons Atomic mass is the average of the isotopes weighted in their abundance Carbon has an atomic mass of 12.01 ATOMS Types of Isotopes Stable Isotopes – their nuclei doesn’t lose subatomic particles (decay) Example: Carbon-12 and Carbon-13 Radioactive Isotopes – nucleus spontaneously decays Gives off particles and energy If decay leads to a change in proton number, it transforms into a different atomic element Example: Carbon-14 decays into a Nitrogen atom ATOMS Applications for Radioactive Isotopes Diagnostic Tool – incorporated into biologically active molecules Used as tracers to track atoms during metabolism Inject small doses into the blood and analyze the tracer molecule secreted in the urine Example: Kidney disorders Example: PET scans to monitor cancer growth ATOMS Applications for Radioactive Isotopes Radiometric (absolute) Dating – measure radioactive decay in fossils to date certain events o Layering of fossil beds can’t determine actual age of fossils based on position alone o Parent isotope decays into daughter isotopes at a fixed rate o Expressed as the half-life of the isotope  The time it takes for 50% of the parent isotope to decay  Each radioactive isotope has a characteristic half-life  Example: Uranium-238 has a half-life of 4.5 ATOMS Energy – capacity to cause change Example: by doing work Varies in the electrons of each atom Only the electrons of each Potential Energy – the energy that atom interact during matter has due to its position or chemical reactions structure Matter tends to move toward the lowest possible state of potential energy Work must be done to restore ATOMS Potential Energy Electrons Created by the distance from the nucleus (position) Opposite attraction to the positively charged protons Total energy is made up of potential energy and kinetic energy When excited, electrons move to a higher energy orbital which creates a higher level of potential energy (requires work) ATOMS Electron Shells – different electron energy levels Represented by concentric circles An electron’s energy level relates to its distance from the nucleus Higher energy level equals a further distance from the nucleus ATOMS Electron Shells – different electron energy levels Represented by concentric circles An electron’s energy level relates to its distance from the nucleus Higher energy level equals a further distance from the nucleus  1st shell: closest to the nucleus, electrons have the lowest potential energy  3rd shell: furthest from the nucleus, electrons have the ATOMS Electron Shells – different electron energy levels Represented by concentric circles An electron’s energy level relates to its distance from the nucleus Higher energy level equals a further distance from the nucleus  Energy absorbed = electron moves to a higher shell Example: light energy is absorbed  Energy is lost = electron moves to a lower shell  ATOMS Distribution of Electrons Determines the chemical behavior of an atom Arranged in three periods on the periodic table o Correspond to the number of electron shells in their atoms Arranged left to right corresponding to the sequential addition of electrons and protons  ATOMS Distribution of Electrons Example: Hydrogen and Helium are both in the first shell Hydrogen has 1 electron Helium has 2 electrons  ATOMS Distribution of Electrons First shell: 2 electron maximum Second shell: 8 electron maximum Third shell: 18 electron maximum  Electrons tend to exist in lowest available ATOMS Distribution of Electrons First shell: 2 electron maximum Second shell: 8 electron maximum Third shell: 18 electron maximum  Electrons tend to exist Example: Lithium has 3 electrons The first two electrons fill the first in lowest available shell states of potential The last electron occupies the energy second shell ATOMS Valence Shell – outermost shell of an atom The electrons present in the outermost shell are called valence electrons The number of valence electrons determines the chemical behavior of the atom Atoms with the same number of valence electrons have similar chemical behavior ATOMS Inert Elements – chemically unreactive Contain completed valence shells (full) Examples: Helium, neon, argon, krypton, xenon, radon ATOMS Electron Orbitals – describes the space where an electron is found 90% of the time More realistic depiction than using two-dimensional, concentric-circle diagrams of electron shells No more than 2 electrons occupy a single orbital First shell: one spherical orbital (1s) Second shell: four orbitals Video Help! One large spherical orbital (2s) ATOMS ATOMS Atomic Interactions Occur between unpaired electrons Goal is to complete their valence shells to become stable CHEMICAL BONDING Chemical Bonds – interactions between atoms to form chemical compounds Share or transfer valence Hydrogen electrons Bond Types of Chemical Bonds Covalent bonds Ionic bonds Hydrogen bonds Van der Waals interactions CHEMICAL BONDING Covalent Bonds – sharing of a pair of valence electrons by two atoms Strongest chemical bond in aqueous environments +2 atoms held together by covalent bonds creates a molecule CHEMICAL BONDING Expressing Covalent Bonds (Example: Hydrogen) Molecular formula – indicates which atoms are present  H2 Lewis dot structure – element symbols are surrounded by dots representing their valence electrons  H:H Structural formula (Lewis structure) – implement lines to depict shared electrons  H—H CHEMICAL BONDING Valence – the bonding capacity of an atom Equals the number of electrons required to complete the valence (outermost) shell Corresponds to the number of covalent bonds the atom can form Determines whether a single bond or double bond results  Single bonds – sharing between one pair of valence electrons  Double bonds – sharing between two pairs of valence electrons CHEMICAL BONDING Pure Elements – composed of only one type of atom Examples: O2, H2 Compounds – formed by +2 different types of elements that are chemically bonded in fixed proportions Examples: H2O, CH4 CHEMICAL BONDING Electronegativity – measure of an atom’s ability to attract shared electrons to itself Stronger the electronegativity = the stronger the atom pulls the shared electrons toward itself Weaker electronegativity = the less the atom pulls the shared electrons toward itself CHEMICAL BONDING Nonpolar Covalent Bonds – when the atoms share electrons equally No electrical charge Two atoms of the same element or have a similar electronegativity Examples: CO2, CH4, Hydrocarbons (gases), diatomic elements (H2, N2, O2, Cl2) CHEMICAL BONDING Polar Covalent Bonds – when the atoms share electrons unequally One atom is more electronegative than the other Degree of polarity varies Creates a partial, electrical charge (𝛅+ or 𝛅-) Examples: H2O, SO2, NH3 CHEMICAL BONDING Ionic Bonds – one atom loses valence electrons and gains them from another atom The more electronegative atom gains an electron from another atom The less electronegative atom loses it electron Results in ions (two, oppositely-charged atoms) CHEMICAL BONDING Cation Loses its electron Positively charged (+) Contain more protons than electrons  t in cation is positive! Anion Gains an electron Negatively charged Contain more electrons than protons  an in anion is a negative! CHEMICAL BONDING Ionic Bond – one atom loses valence electrons and gains them from another atom Attraction between opposite charges o Example: sodium (11Na) and chlorine (17Cl) o Lone sodium valence electron is transferred to the chlorine atom o Both atoms complete their valence shells o Sodium loses its electron and becomes a cation (+1 charge) CHEMICAL BONDING Ionic Compounds (salts) – compounds formed from ionic bonds Strongest bonds in dry environments Examples: Sodium chloride (NaCl), Sodium hydroxide (NaOH), Lithium chloride (LiCl) CHEMICAL BONDING Hydrogen Bonds – noncovalent attraction between a hydrogen atom and an electronegative atom Results from a hydrogen atom covalently bonding to an electronegative atom Creates a partial charge and is attracted to another electronegative atom nearby Weak chemical interaction Examples: water, ammonia, hydrogen fluoride CHEMICAL BONDING Van der Waals Interactions – very weak attraction between molecules that are very close to each other Unequally distributed positively and negatively charged regions Individually very weak CHEMICAL BONDING Van der Waals Interactions – very weak attraction between molecules that are very close to each other Unequally distributed positively and negatively charged regions Individually very weak CHEMICAL REACTIONS Chemical Reactions – changing the composition of matter by making and breaking chemical bonds  Reactants – starting materials  Products – resulting materials Other Information: Arrow – indicate the conversion of reactants to products Coefficients – indicate the number of those molecules involved All reactant atoms must be accounted for in the products!  The Law of Conservation of Mass