The Building Blocks of Matter: A Chemistry Primer

Chemistry Elements = a substance that can’t be broken down into simpler substances by chemical or physical means • 92 elements occur naturally on Earth • H & He – the most abundant elements in the universe o make up ~ 99.8% of the universe (by mass) • C, H, O, N – make up > 96% of the mass of a human • O, Si, Al, Fe, Ca, Na, K, Mg – make up > 98% of the Earth’s crust o O & Si – make up ~ 75% of the crust (most of the rocks and minerals) Chemical Symbol = one or two letter abbreviation for an element H = hydrogen; He = helium; C = carbon; O = oxygen; etc. Atoms = the smallest particle of an element that has the characteristics of that element Nucleus – made up of protons (has a positive electrical charge: p+) & neutrons (has a neutral electrical charge: n0) • Atomic # = # of protons; determines what element it is • Mass # = ave. mass of element and all isotopes • Strong Force = the energy that binds the nucleus together Ex. 9 F Fluorine 18.998 Chemical Symbol = F Atomic # = 9 Mass # = 19 Element = fluorine # Protons = 9 # Neutrons = 19 – 9 = 10 Electron Cloud – electrons surround the nucleus of an atom (they have a negative electrical charge: e–) • Unlike protons and neutrons, electrons have relatively no mass • Electron charges have the same magnitude as proton charges (so, in a stable atom, the number of electrons will equal the number of protons and their charges will cancel out) • Energy Levels – represent the area in an atom (around a nucleus) where an electron is most likely to be found o The first energy level can hold up to 2 electrons o The second level hold 8 electrons (3rd holds 18, etc.) o Valence Electrons = the electrons in the outermost energy level  They determine the chemical behavior of the different elements Isotopes = atoms of the same element that have different numbers of neutrons (12C vs. 14C) • radioisotopes & radiometric dating • Radioactivity = the spontaneous process through which unstable nuclei emit radiation o During radioactive decay, the nucleus can lose protons and/or neutrons, change a proton to a neutron, or change a neutron to a proton o Because the number of protons identifies the element, decay can change the identity of that element Compounds = a substance composed of atoms of two or more elements that are chemically combined (e.g. H2O or NaCl) • the proportions of individual atoms are fixed • form depending on # and arrangement of electrons in the atoms • properties of compounds are different than that of the individual parts – ex. NaCl • atoms tend to combine to form compounds that lend stability = octet rule – want to fill valence shell (8 outermost electrons) Chemical Bonds = the forces that hold the elements together in a compound - generally covalent, ionic, or metallic bonds (you’ll learn more in chemistry) Chemical Formula = the type element and number of atoms in a compound • H2SO4 & Mg3(PO4)2 : subscript represents the # of atoms Covalent Bonding = sharing of electrons causes an attraction between 2 atoms • tend to be 2 non-metals Molecule = a group of atoms held together by covalent bonds and having no overall charge (because the # of electrons equals the # of protons) Polar Molecules = an unequal sharing of electrons by the atoms in a molecule • the oxygen atom, in a water molecule, does not share the electrons equally with the 2 hydrogen atoms • this “unequal sharing” causes the molecule to be “partially” charged and causes them to be attracted more strongly than they would be by simple covalent bonding • has a positive end and a negative end Ions = A charged particle that has either gained or lost electrons • an atom that has 4 or more valence electrons tends to fill its outer shell by adding electrons o they will form negative ions • an atom that has less than 4 valence electrons tends to fill its outer shell by losing electrons o they will form positive ions Ionic Bonding = the attractive force b/w 2 ions of opposite charge - gain/loss of electrons b/w atoms; form ions in solution • tend to be a metal and non-metal • superscript represents the charge of an ion (ion = charged particle) • compounds formed by ionic bonding are known as ionic compounds Metallic Bonds • in metals, the valence electrons are shared by all the atoms • the positive ions of the metal are held together by the negative electrons between them • this allows metals to conduct electricity (because the electrons can move freely throughout the entire solid metal) Chemical Reactions – when chemical reactions occur, bonds b/w atoms are formed or broken, causing substances to combine and recombine as different molecules • 2H2 + O2  2H2O (hydrogen + oxygen  water) 4Fe + 3O2  2Fe2O3 (iron + oxygen  hematite) reaction is represented by chemical equations Mixtures = a combination of substances in which the individual components retain their original chemical properties (stir sand & sugar) • heterogeneous (can see the different parts; such as soil) • easily separated by physical means Solutions = a mixture in which one or more substances are uniformly/evenly distributed in another substance = a homogeneous mixture (can’t distinguish parts; uniform throughout) - one substance (solute) is dissolved into another (solvent) and does not settle out of solution Solute = the substance being dissolved in a solution - always the lesser volume of the two Solvent = the substance doing the dissolving - always the greater volume of the two Concentration = a measurement of the amount of solute dissolved in a fixed amount of solution (the more solute dissolved = the greater the concentration of the solution) Acids = a solution containing a substance that produces hydrogen ions (H+) in water H2O + CO2  H2CO3 - (carbonic acid) some of this acid dissociates (breaks apart) in water into ions H2CO3  H+ + HCO3– • • (hydrogen ion & bicarbonate ions) have a pH below 7 tend to have a sour taste, produce tingly or burning sensation on the skin, and are highly corrosive in concentrated form Bases = produce hydroxide ions (OH–)in solution • bases can neutralize acids (to some degree) by combining with hydrogen ions to form water H+ + OH–  H2O • • have a pH above 7 tend to have a bitter taste and feel slippery pH = the measure of how acidic or basic (alkaline) a solution is • measured on a scale of 0 – 14, with 7 being neutral States of Matter All molecules have molecular motion; depending on the amount of motion (and therefore energy) that molecule has, is what determines the state of that matter • Molecular motion/vibrations increase with temperature (so the particles move faster and have more energy as they change from a solid to a liquid to a gas) • All matter has the ability to take any state (you just need to add or take away enough energy – usually in the form of heat) Shape Volume Energy Solids Liquids Gases Plasma Solids • • Liquids • Gases • • • • • Plasma • • • Most solids are densely packed particles Most solids have a crystalline structure o Crystals form symmetrical solid objects with flat faces and straight edges between the faces At the melting point of matter (which is different for each substance), the vibrations of the matter become sufficient enough to break the forces holding the solid together The particles can then slide past each other; thus becoming a liquid When particles in a liquid gain enough energy they can escape to become a gas This is called evaporation, or vaporization When a solid goes directly from a solid to a gas, skipping the liquid state, it is called sublimation (examples are moth balls or ice in the freezer) When a gas is cooled, it releases thermal energy in a process called condensation, changing into a liquid When matter is heated to temperatures greater than 5000°C, the collisions between the particles are so great that electrons are knocked away from atoms Such extreme temperatures exist in stars o As a result, the gases of stars consist entirely of positive ions and free electrons These hot, highly ionized, electrically conducting gases are called plasmas Law of Conservation of Matter & Energy • Matter can neither be created or destroyed (only changed from one state to another) • Energy cannot be created or destroyed (only changed from one form to another)