Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
CHEMISTRY: ATOMS, MOLECULES & REACTIONS HAGEBUSCH NEO A&M – PHYS 1014 Atomic Theory Atom – smallest part of an element that still retains the properties of that element Atomic Theory Dalton’s Atomic Theory: 1. All elements are composed of atoms 2. All atoms of an element are identical and differ from atoms of other elements 3. Atoms of different elements can chemically combine in simple whole number ratios 4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms are never turned into a different atom during a reaction Atomic Structure Electron (e-) – subatomic particle with a negative charge • Discovered by English physicist JJ Thompson, 1897 • Smallest of all subatomic particles • Electrons are located in empty space around the nucleus called an electron cloud. The Electron Streams of negatively charged particles were found to emanate from cathode tubes. The Atom, circa 1900 The prevailing theory was that of the “plum pudding” model, put forward by Thompson. It featured a positive sphere of matter with negative electrons imbedded in it. Rutherford’s Gold Foil Experiment Rutherford’s Gold Foil Experiment Atomic Structure Atomic Nuclei - central part of an atom that contains about 99.9% of atomic mass 1. Proton (p+) – subatomic particle with a positive charge Discovered by Rutherford, 1919 2. Neutron (no) – subatomic particle with no charge Discovered by James Chadwick, 1932 Atomic Structure The Bohr Model of the Atom • Neils Bohr, a Danish physicist, proposed that electrons occupy specific orbits around the nucleus called energy levels • Energy level – region around the nucleus of an atom where an electron is likely to be moving Atomic Structure The Bohr Model of the Atom He proposed that the electrons have a fixed energy and that keeps them away from the nucleus The closer the energy level is to the nucleus, the lower the energy of the electrons in it Bohr discovered this by studying the emission spectrum emitted by atoms after they were charged with high voltage electricity Subatomic Particles Protons and electrons are the only particles that have a charge. Protons and neutrons have essentially the same mass. The mass of an electron is so small we ignore it. Distinguishing Between Atoms Atomic Number – the number of protons in the nucleus of an atom • • • • Always a whole number on the periodic table In a neutral atom, the number of protons and electrons is equal The number of electrons can change and the element remains the same If the number of protons changes, the element changes Distinguishing Between Atoms Mass Number – total number of protons and neutrons in an atom • Measured in atomic mass units (amu) – 1/12 of the mass of a carbon atom • Always a decimal number • Number of neutrons = mass # - atomic # • Atomic mass is actually the average of an atom of that particular element • Isotopes – atoms with the same number of protons, but a different number of neutrons Isotopes Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons. 11 C 6 12 C 6 13 C 6 14 C 6 Using the Periodic Table Atomic # C 6 Symbol 12.011 Mass # Periodic Table Development of the Periodic Table • • Early scientists grouped elements according to similar properties Dimitri Mendeleev, a Russian chemist, was the first to organize the all elements in a systematic way o o At first there were only 70 elements Mendeleev even left spaces where he knew would be discovered Periodic Table The Modern Periodic Table 112+ elements on the periodic table Horizontal rows are called periods. Vertical columns are called groups or families • • • o o All elements in a group have similar physical and or chemical properties because they have the same number of valence electrons Valance Electron – electrons in the outermost energy level of an atom V. Periodic Table The Modern Periodic Table • Arranged by increasing atomic number • Periodic Law (periodicity) – pattern on the periodic table where certain properties repeat when the elements are arranged according to atomic number Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities. METALS Solids at room temp • High electrical conductivity • • Lustrous • Ductile • Malleable Non-metals Various states of matter at room temp • Poor electrical conductors • • Non-Lustrous • Brittle if solid Metalloids Exhibit properties of both metals and nonmetals • Periodicity Main Group Elements ( Group 1, 2, 13 – 18 ) • • • • Alkali Metals (Group 1) – most highly reactive main group metals on the periodic table Alkali metals are so reactive they are not found in pure form on earth Alkaline Earth Metals (Group 2) – exhibit similar properties to group 1, but are slightly less reactive Halogens (Group 17) – most highly reactive group of non-metals Noble Gases (Group 18) – also called inert gases, least reactive group of elements on the periodic table and undergo very few chemical reactions Periodicity Transition Metals (Group 3 – 12) – largest group of elements on the periodic table and contains about 80% of the metals found on Earth Inner Transition Metals – bottom two rows of elements on the periodic table and the least common elements found on Earth Periodicity Inner-transition Metals • Lanthanide Series – naturally occurring elements with very similar properties • Actinide Series – contain most of the radioactive elements and all have similar properties The Periodic Table INNER-TRANSITION METALS H Actinide Series Noble Gases Lanthanide Series Halogens Alkaline Earth Metals Alkali Metals Transition Metals Chemical Bonding Bonding occurs between atoms in order to redistribute their valence electrons in a more stable way. The manner of redistribution determines the type of bonding that takes place Chemical Bonding Types of Chemical Bonds 1. 2. Ionic bonding – chemical bonding resulting from the electrical attraction between cations and anions Covalent bonding – results from the sharing of valence electrons between two atoms Chemical Bonding Factors that effect bonding 1. Electronegativity – measure of an atoms ability to attract electrons from another atom in a chemical bond Electronegativity is measured from 4.0 to 0, the higher the number, the higher the electron affinity or ease of gaining eElectronegativity and the periodic table Alkali & Alkaline Earth metals are the least electronegative Non-metals are more electronegative Fluorine is the most electronegative Electronegativity Electronegativity Chemical Bonding Factors that effect bonding 2. Electron structure An atom’s electron configuration is the key factor in determining its electronegativity Electron-dot notation – visual representation of an atoms valence electrons Each atom strives to obtain the same valence electron structure as a noble gas Each noble gas has a full valence energy level because the s & p sublevels are full Octet rule – atoms combine in such a way as to completely fill their highest occupied energy level Chemical Bonding Factors that effect bonding 2. Electron structure An atom’s electron configuration is the key factor in determining its electronegativity Electron-dot notation – visual representation of an atoms valence electrons Each atom strives to obtain the same valence electron structure as a noble gas Each noble gas has a full valence energy level because the s & p sublevels are full Octet rule – atoms combine in such a way as to completely fill their highest occupied energy level Chemical Bonding Results of chemical bonding When atoms react and combine with one another they form chemical compounds Chemical compounds are represented by chemical formulas Chemical formulas – show the relative numbers of atoms of each kind in a compound by using atomic symbols Ex: H2O CO2 C6H12O6 NaCl CaBr2 Chemical Bonding Types of Chemical Compounds 1. 2. Covalent Compound – group of atoms held together by the sharing of valence electrons Ionic Compound – group of atoms held together through the attraction of oppositely charged ions Chemical Bonding Covalent compounds are always composed of one or more non-metals Covalent compounds form when atoms share their valence electrons Each atom shares their electrons with the purpose to have eight valence electrons, or in accordance with the Octet Rule Covalent Compounds Exceptions to the octet rule Hydrogen – forms only one bond for a total of only two valence electrons Boron – forms one bond with each of it’s valance electrons for a total of six valence electrons Other elements, such as Phosphorus, can have an expanded valence level if highly electronegative elements are involved Ex: PF5 or SF6 Covalent Compounds Types of Covalent Compounds 1. Polar-covalent compound – covalent compound in which there is an unequal distribution of electrons 2. Polar compounds result when an atom in a compound does not share the electrons equally across a bond Nonpolar-covalent compound – covalent compound where all electrons are shared equally Covalent Compounds Characteristics of Covalent Bonds Bond Length – distance needed between covalently bonded atoms to minimize their potential energy The bond length is different depending on the atoms involved The bond length is ideal when the distance between the atoms allows the repulsion of like charges to equal the attraction of opposite charges Covalent Compounds Characteristics of Covalent Bonds Bond Energy – energy needed to break a covalent bond The potential energy of the original atoms is stored in the bonds Measured in kilojoules per mol (kJ/mol) Covalent Compounds Characteristics of Covalent Bonds Bond Types Single Bond – covalent bond in which only two electrons are shared between atoms Multiple Bonds – covalent bond in which multiple pairs of electrons are shared between atoms o o Double Bond Triple Bond Covalent Compounds Characteristics of Covalent Bonds Lewis Structure – using symbols, lines, and dots to represent the structure and shape of a molecule LONE PAIR Structural formulas are used to indicate the basic shape of a molecule without the use of dots for the lone pairs O FREE e- Covalent Compounds Characteristics of Covalent Bonds Resonance Structures – refers to compounds that cannot be represented following general Lewis structure procedures Resonance structures occur when some elements both share and transfer electrons to achieve the octet rule Ex: Ozone ( O3 ) Carbon Monoxide ( CO ) Sulfur Dioxide ( SO2 ) Ionic Compounds Ionic compounds usually contain at least one metal Ionic compounds are composed of atoms held together due to the attraction of oppositely charged ions Ion – any charged particle Cation – positively charged ion Created when an atom loses electrons Anion – negatively charged ion Created when an atom gains electrons Ionic Compounds Characteristics of Ionic Bonds Ionic compounds arrange themselves in a crystal lattice Ionic Compounds Characteristics of Ionic Bonds Ionic compounds arrange themselves in a crystal lattice Lattice Energy – energy released when one mole of an ionic crystalline compound is formed from gaseous ions Positive energy values indicate energy is absorbed and stored in the bond Negative energy values indicate energy is released Compound Comparison Ionic Cmpds Contains ions Usually contains a metal Strong attraction between formula units High melting and boiling point Conduct electricity when dissolved in water Covalent Cmpds valence e Contains nonmetals Weak atraction between molecules Low melting & boiling points Shares Polyatomic Ions Polyatomic ion – charged group of covalently bonded atoms Charge results when there is either too many or too few electrons to adequately satisfy the octet rule for each of the atoms involved Polyatomic ions bond with oppositely charged ions to form part of an ionic bond Chemical Formulas The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound. © 2009, Prentice-Hall, Inc. Names and Formulas Chemical formula – indicates the types and the number of atoms in a compound Covalent compound = molecule Ionic compound = formula units Chemical formulas serve as a rough blueprint for a molecule * First Step * Determine if the compound is IONIC or COVALENT MgBr2 1. Ionic metals / Covalent non-metals Names and Formulas Ionic Compounds 2. Determine if the compound is Binary or Tertiary • • Binary compound – composed of only two different elements Tertiary compound – composed of three or more different elements • All tertiary ionic compounds contain at least one polyatomic ion MgBr2 1. Ionic metals 2. Binary / Tertiary 2 elements 3+ elements / Covalent non-metals Names and Formulas Binary Ionic Cmpds w/ Main Group Metals 3. Determine if the cation is a Main Group Metal • Group 1, 2 or Aluminum MgBr2 1. Ionic metals / Covalent non-metals 2. Binary / Tertiary 2 elements 3+ elements 3. Main Group / Transition Metal Group 1, 2, Al Names and Formulas Binary Ionic Cmpds w/ Main Group Metals 3. Determine if the cation is a Main Group Metal • Group 1, 2 or Aluminum • Name the cation • Name the anion and change the ending to -ide MgBr2 1. Ionic metals / Covalent non-metals 2. Binary / Tertiary 2 elements 3+ elements 3. Main Group / Transition Metal Group 1, 2, Al A. Name cation (metal) B. Name anion & change ending to “-ide” Magnesium Bromide Li2CO3 1. Ionic metals 2. Binary / Tertiary 2 elements 3+ elements / Covalent non-metals Names and Formulas Tertiary Ionic Cmpds Tertiary ionic compounds contain polyatomic ions Use the chart on p 226 to identify the polyatomic ion Use basically the same rules for naming binary ionic compounds, but DO NOT change the ending of the polyatomic ion Li2CO3 1. Ionic metals 2. Binary / Tertiary 2 elements 3+ elements / Covalent non-metals A. Must include polyatomic ion B. ID the polyatomic ion, use same rules for binary but DO NOT change the ending 3. Main Group / Transition Metal Group 1, 2, Al A. Name cation (metal) B. Name anion & change ending to “-ide” (No change for polyatomic ion) Lithium Carbonate Fe2O3 1. Ionic metals / Covalent non-metals 2. Binary / Tertiary 2 elements 3+ elements 3. Main Group / Transition Metal Group 1, 2, Al Any other metal Names and Formulas Ionic Cmpds w/ Transition Metals Transition metals are different because the have multiple oxidation numbers Oxidation # - charge associated with an ion in an ionic compound Determined based on the amount of electrons lost or gained Main group elements and polyatomic ions have constant oxidation numbers Oxidation numbers must be calculated for the transition metals according to the anion it is bonded to Calculating Oxidation #’s 1. 2. 3. Determine the ox # for the anion (non-metal or polyatomic ion) Calculate the total charge contributed by the anion Calculate the unknown ox # for the transition metal by balancing the charges Ion Ox # X How many in cmpd Total = Charge Names and Formulas Covalent Compounds Covalent compounds are composed of nonmetals bonded together covalently Covalent compounds do not rely on charge balance Naming of covalent compounds uses prefixes to indicate the number of each element in the compound Many covalent compounds have “names” instead of the stock system of naming Ex: C6H12O6 – glucose CH4 - methane Names and Formulas Naming Covalent Compounds Name the compound using the prefixes Mono is understood if the first element only has one atom ALWAYS include a prefix on the second element # Prefix # Prefix 1 2 3 4 5 mono 6 7 8 9 10 hexa di tri tetra penta hepta octa nona deca Describing Chemical Reactions Chemical reaction – process where one or more substances are changed into one or more different substances with new chemical properties Chemical equation – uses symbols, formulas and molar ratios of compounds to describe a chemical reaction 6CO2 + 6H2O → C6H12O6 + 6O2 reactants products Indications of a Chemical Reaction 1. 2. 3. 4. Production of light or a transfer of heat/NRG Production of a gas Formation of a precipitate (ppt) Color change Characteristics of a Chemical Equation 1. The equation must represent known facts 2. The equation must contain correct formulas for reactants and products 3. Must adhere to the Law of Conservation of Mass Matter cannot be created or destroyed All atoms on the reactant side of the reaction must equal the atoms on the product side Balancing Chemical Equations CH4 + O2 → CO2 + H2O Make a list of each element on both sides Start adding COEFFICIENTS to the compounds to balance the atoms on each side Balance O’s, H’s, and other single elements last 5 Types of Chemical Reactions 72 1. Synthesis reaction (composition reaction) – two or more substances combine to form a new, more complex compound 2. A + B → AB Decomposition reaction – a single compound undergoes a reaction in which two or more simpler substances are produced AX → A + X NRG Hagebusch - Chemistry I 4/9/2012 5 Types of Chemical Reactions 73 1. Single-Displacement reaction (replacement reaction) – one element replaces a similar element in a compound 1. A + BX → AX + B Double-Displacement reaction – the ions of two different compounds exchange places in an aqueous solution to form two new compounds AX + BY → AY + BX Hagebusch - Chemistry I 4/9/2012 5 Types of Chemical Reactions 74 5. Combustion reaction – a substance combines with oxygen, releasing a large amount of energy CH4 + O2 → CO2 + H2O Hagebusch - Chemistry I 4/9/2012 Synthesis Reactions 75 RXNS of Elements with Oxygen & Sulfur Most all metals will combine with O & S to form Oxides & Sulfides Oxygen and Sulfur are both Group 16 elements and have oxidation #’s of -2 Nonmetals also under go reactions with Oxygen to produce oxides 2Mg(s) + O2(g) → 2MgO 2H2(g) + O2(g) → 2H2O(g) Hagebusch - Chemistry I 4/9/2012 Synthesis Reactions (cont.) 76 RXNS of Metals with Halogens Most metals react with Group 17 elements to form metallic salts Fluorine is the most reactive element and will undergo synthesis with almost every element 2Na(s) + Cl2(g) → 2NaCl(s) Mg(s) + F2(g) → MgF2(s) Hagebusch - Chemistry I 4/9/2012 Synthesis Reactions (cont.) 77 RXNS of Metals with Oxides The oxides of active metals, such as group 1 & 2, react with water to form metal hydroxides CaO(s) + H2O(g) → Ca(OH)2(s) Mg(s) + F2(g) → MgF2(s) Oxides of nonmetals react with water to form oxyacids SO2(g) + H2O(l) → H2SO3(aq) Hagebusch - Chemistry I 4/9/2012 Decomposition Reactions 78 Decomp. of Binary Compounds Simplest form of decomposition Electrolysis – decomposition of a compound by electric current H2O(l) → 2H2(g) ELECTRICITY + 2O2(g) Binary oxides decompose through the addition of heat Δ HgO(s) → 2Hg(s) + O2(g) Hagebusch - Chemistry I 4/9/2012 Decomposition Reactions (cont.) 79 Decomp. of Metal Carbonates Products are a metal oxide and CO2(g) Δ CaCO3(s) → CaO(s) + CO2(g) Decomp. of Metal Hydroxides Products are a metal oxide and H2O(l) Δ Ca(OH)2(s) → CaO(s) + H2O(l) Decomp. of metal Chlorates Products are a metallic salt and O2(g) Δ 2KClO3(s) → 2KCl(s) + 3O2(g) Hagebusch - Chemistry I 4/9/2012 Decomposition Reactions (cont.) 80 Decomp. of Acids Certain acids decompose to form non- metal oxides and H2O(l) H2CO3(aq) → CO2(g) + H2O(l) Hagebusch - Chemistry I 4/9/2012 Single-Displacement Reactions 81 Metal Replacement Occurs when a more reactive metal replaces a less active metal in a compound 2Al(s) + 3Pb(NO3)2(aq) → 3Pb(s) + 2Al(NO3)3(aq) Displacement of H+ in H2O by a Metal Active metals, such as Group 1 react readily with water to produce metal hydroxides and hydrogen gas 2Na(s) + 2H2O(l) → 2NaOH(aq) Hagebusch - Chemistry I 4/9/2012 + H2(g) 82 Single-Displacement Reactions (cont.) Displacement of H+ in an Acid by a Metal Some active metals will replace H+ in acids Mg(s) + 2HCl(aq) → MgCl2(s) + H2(g) Displacement of Halogens One halogen replaces another in a compound Fluorine is the most reactive Cl2(g) + 2KBr2(aq) → KCL2(aq) Hagebusch - Chemistry I + 4/9/2012 Br2(l) Double-Displacement Reactions 83 Formation of a Precipitate (ppt) A precipitate is formed when the cation from one compound reacts with the anion from another to form an insoluble compound 2KI(aq) + Pb(NO3)2(aq) → 3PbI2(s) + 2KNO3(aq) Formation of a Gas FeS(s) + 2HCl(aq) → 2H2S(g) + FeCl2(aq) Hagebusch - Chemistry I 4/9/2012 84 Double-Displacement Reactions (cont.) Neutralization When strong acids are mixed with strong bases, the product is a metallic salt and water HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(aq) Hagebusch - Chemistry I 4/9/2012 Combustion Reactions 85 Spontaneous reaction between a compound and oxygen gas, producing water vapor When organic compounds combust, the products are carbon dioxide and water vapor 2H2(g) + O2(g) → H2O(l) CH4(g) + O2(g) → 2CO2(g) Hagebusch - Chemistry I + H2O(l) 4/9/2012