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Chapter 2- Atoms, Molecules, and Ions CHEMISTRY II Materials- seemingly infinite variety of types and properties Properties can be classified- much harder to understand and explain Use knowledge of how atoms are structured and how they combine with one another Atoms of different elements act differently- we will look at why Sec 2.1- Atomic Theory of Matter Atomic theory has been developed, modified, and revised since 460 BC Democritus- gave atom the name “atomos”meaning “indivisible or uncuttable” Idea of elements were linked with idea of atoms once chemists were able to better measure the amounts of elements that reacted with one another to form new substances Dalton’s Atomic Theory John Dalton developed a set of postulates to explain atomic theory: 1: Each element is composed of extremely small particles called atoms 2: All atoms of a given element have identical mass and properties, but atoms of one element are different from atoms of all other elements 3: The atoms of one element cannot be changed into atoms of another element during a chemical reaction; atoms are neither created nor destroyed in chemical reactions 4: Compounds are formed when atoms of more than on element combine; a given compound always has the same relative number and kind of atoms In Dalton’s Atomic Theory- atoms defined as the smallest particles of an element that retain the chemical identity of the element (unlike Democritus) Element- One type of atom Compound- Combo of atoms of two or more elements Laws of chemical combination Dalton theory: applied several laws to develop his atomic Law of constant composition- in a specific compound, the relative number and kinds of atoms are constant (4th postulate) Law of Conservation of Mass- total mass of materials after a chemical reaction is the same as the total mass before the chemical reaction (3rd postulate) Law of Multiple Proportions Dalton developed Law of Multiple proportions- if 2 elements A and B combine to form more than one compound, the masses of B that can combine with A are in the ratio of small whole numbers Ex: H and O can combine to form different compounds- H2O and H2O2 For H2O- 1.0g H combine with 8.0g O For H2O2- 1.0g H combine with 16.0g O Ratio of mass of O per gram of H between 2 compounds is 1:2 (small whole numbers) We can use this ratio to conclude that H2O2 has twice as many atoms of O per atom of H than H2O does C and O can combine to form different compounds One compound contains 1.33g O for every 1g of C Another compound contains 2.66g O for every 1g of C If the first compound has equal number of O and C atoms, what can we conclude about the composition of the second compound? Sec 2.2- Discovery of Atomic Theory Dalton actually never had hard physical evidence of the existence of atoms Modern instruments have allowed us to gain information on atoms- individual properties and even images of atomsLet’s take a look at the world’s “smallest” movie… Even before the imaging of atoms, we had found evidence of their existence The further we explored atoms, the more we realized that maybe they aren’t indivisible (like Democritus proposed)- they indeed have a more complex structure We found evidence of subatomic particles The current atomic model has been developed thanks to a few landmark discoveries- experiments that led to the discovery of electrons, the nucleus, and protons The electron Cathode ray tubes- tubes with almost no air and electrodes at either end When a voltage is applied to the electrodes, radiation was produced that flowed from the negative electrode (cathode) to the positive electrode (anode) Cathode ray- radiation coming from the cathode Thompson proposed that these cathode rays were actually streams of negatively charged particles, that he named electrons Through experimentation, he was able to measure the amount of electric charge in a given mass of electronsrelationship between charge and mass From Plum Pudding to Nucleus Thompson- Proposed that atoms were positive spheres with embedded electrons (psh) Rutherford- shot some particles at thin gold foilfound that a lot of the shot particles passed through- but more importantly that some of them were being deflected at large angles This gave evidence of areas of concentrated mass inside atoms Rutherford called this area of concentrated mass a nucleus- and proposed that it was made up of a positive charge Rutherford proposed that most of the volume of atom is empty space, in which electrons use to move around Further experimentation lead to the discovery of protons and neutrons, and that these subatomic particles make up the nucleus together Sec 2.3- The Modern View of Atomic Structure List of particles that make up a nuclei is growing longer and longer We will only be concerned with 3 subatomic particles- protons, neutrons, and electrons, since these are the only ones that are associated with chemical behavior Electronic charge Charge of electron= -1.602 x 10-19 C Charge of proton= + 1.602 x 10-19 C Electronic charge= 1.602 x 10-19 C For convenience, we will express charges of protons/electrons in terms of multiples of electronic charge So we will consider a proton to have a charge of 1+, and an electron to have a charge of 1- Every atom has equal numbers of electrons and protons, so atoms themselves have no electronic charge (talking about atoms, not ions) Protons nucleus and neutrons- packed together in the Nucleus- very small compared to the total volume of an atom- densest part of any atom Majority of volume of atoms is space outside of nucleus that electrons occupy- mostly empty Negative electrons attracted to positive protons in nucleus by electrostatic forces- the strength of this attraction can be used to explain differences between different atoms We use atomic mass units (amu) to express the mass of an atom due to the fact that atoms have incredibly small masses (on the order of 10-22 g) Protons and neutrons have similar masses Mass of electron- very small compared to mass of proton/neutron Due to comparatively small mass, electrons do not contribute much to the overall mass of an atom- most mass comes from protons and neutrons (nucleus) Atoms of different elements differ significantly in their subatomic compositions Atoms of each element have a characteristic number of protons Atomic number- the number of protons in the nucleus of an atom of a particular element Atoms, due to net zero charge, have same number of electrons as protons Atom of C has 6 protons (atomic number) and 6 electrons Atom of O has 8 protons and 8 electrons Atoms of a given element may differ in number of neutrons Mass number- Total number of protons and neutrons in a given atom (NOT atomic mass) Mass number will sometimes be written along with an atomic symbol, or tagged on at the end of an element’s written name, to specify what isotope is being dealt with Atoms with identical atomic numbers (protons), but different mass numbers are called isotopes of each other (isotopesdifferent mass number due to different number of neutrons) Only use the notation with superscripts only when referring to a particular isotope Sec 2.4- Atomic Weights Atoms indeed have a mass- we talked about how they are small pieces (the building blocks) of matter Atoms of different elements have different masses Back in the day, atomic masses were originally given to elements on a basis of Hydrogen (H was given the atomic mass of 1, arbitrarily) Now we have instruments that we can use to find the mass of atoms As mentioned earlier, due to the very small value of the mass of an atom in grams, we conveniently use atomic mass units (amu) to talk about atomic mass Amu is based off of the carbon-12 atom. (12C= 12 amu, exactly) Most elements exist in nature as a mixture of isotopes Average atomic mass- calculated using the mass of an elements’ various isotopes and their relative abundances Average atomic mass is also known as atomic weightthis is the number (often a decimal) that you see listed under an element’s symbol on the periodic table As stated earlier- we can calculate an element’s average atomic mass (atomic weight) using the mass of various isotopes and their known abundances. Ex: Chlorine naturally occurs as two isotopes: 35Cl (atomic mass=34.969 amu) and 37Cl (atomic mass= 36.966 amu) 35Cl is found in 75.78% abundance, and 37Cl is found in 24.22% abundance Calculate the average atomic mass (atomic weight) of Chlorine Ever wonder how in the heck we know how “abundant” each isotope of each element is? Mass Let’s Spectrometry- Lovingly called “Mass spec” take a look at how mass spec works… (courtesy of Bozeman Science) Sec. 2.5- The Periodic Table Sec. 2.5- The Periodic Table After Dalton’s atomic theory, chemical experimentation became somewhat of a fad Body of chemical observation grew, list of known elements expanded, people tried to find a pattern in chemical behavior between different elements These attempts in identifying patterns resulted in the formation/organization of elements into the periodic table “Periodic” - refers to the patterns seen in chemical behavior (they are predictable and repeat themselves) The periodic table is the most important tool chemists use to organize and remember chemical facts When you arrange elements in order of increasing atomic number, you will see a repeating (or periodic) pattern of chemical/physical properties Periodic table- elements arranged in order of increasing atomic number, with elements of similar properties placed in vertical columns Each element typically has its atomic symbol, atomic number, and atomic weight Vertical columns of periodic table= groups (or families) Horizontal rows of periodic table= periods Elements in the periodic table can be broken into metals, nonmetals, and metalloids Metals- Make up most of the periodic table All elements (except H) to the left of the metalloid staircase are metals All elements (plus H) to the right of the metalloid staircase are nonmetals Metalloids- elements touching the “staircase” that separates the metals from the nonmetals Metals- in general, all share characteristic properties such as luster and high conductivity of electricity and heat. All except Mercury are solid at room temp Nonmetals- in general, differ from metals in appearance and physical properties- nonmetals don’t conduct heat/electricity well, and are also a mixture of solid, liquid, and gases at room temp Metalloids- Share properties of both metals and nonmetals Common names of Groups Group 1A- Alkali metals Group 2A- Alkaline earth metals Group 7A- Halogens Group 8A- Noble gases Sec 2.6- Molecules and molecular compounds Most atoms are actually not found in nature by themselves- this is due to reactivity What group of elements is the most stable? Noble gases- only elements normally found as isolated elements in nature Most matter found in naturemolecules or ions (formed from atoms) Molecule- assembly of two or more atoms bonded together Behaves Think in many ways as a single blob of any object that has many components, but is recognized as a single object (car, phone, etc) Many elements found in molecular form- two or more of the same type of atom (element) are bound together Naturally occurring Oxygen exists as a molecule, with a chemical formula of O2 The “2” subscript means there are 2 Oxygen atoms in every 1 molecule of Oxygen Molecules that consist of two atoms are called diatomic molecules Naturally occurring diatomic moleculesHOBrFINCl (H2, O2, Br2, F2, I2, N2, and Cl2) Molecular compounds- compounds that are composed of molecules and contain more than one type of atom Composition of each compound given by its chemical formula H2O- one molecule consists of 2 H and 1 O (subscripts of chemical formula tell us how many atoms of each element are in one molecule) Most molecular substances we will encounter will contain only nonmetals Molecular vs Empirical formula Molecular formula- chemical formulas that indicate the actual numbers and types of atoms in a molecule Empirical formula- Chemical formulas that give only the relative number of atoms of each type in a molecule Subscripts in an empirical formula- always smallest possible whole-number ratios Ex: H2O2 (molecular formula) simplifies to empirical formula of HO Some substances have the same molecular and empirical formulas Ex: H2O, CO2 Molecular formulas give us more information about molecules than empirical formulas do- they tell us the exact number of each atom in the molecule Molecular formula can give us empirical formula, but having just an empirical formula can not possibly give us the molecular formula (we would need empirical plus more info) Why use empirical formulas? Some methods of sample analysis will only give the empirical formula Once empirical formula is known- additional experiments can give us information needed to convert the empirical formula to the molecular formula Picturing Molecules Molecular formula may tell us the composition of a molecule, but it does not tell us how the atoms are arranged (structure) Structural formula- shows which atoms are attached to which within a molecule Atoms represented with chemical symbols, lines represent bonds between atoms Geometry is not depicted in structural formulas- its’s hard to picture geometry on a flat surface Various models used to visualize molecular geometry- Ball-and-stick models show atoms as spheres (balls) and bonds as sticks Ball-and-stick- can accurately represent angles at which atoms are bonded within the molecule Space-filling model- depicts what the molecule would look like if the atoms were scaled up in size Ball-and-stick Space-filling Sec. 2.7- Ions and Compounds If electrons are gained or lost by a neutral atoms, a charged particle called an ion is formed Cation- ion with positive charge Anion- ion with negative charge Net charge of an atom is represented by a superscript Net positive- comes from losing electrons Net negative- comes from gaining electrons Polyatomic ions- atoms joined together as molecules, but together they have a net negative or net positive charge Ex: NH4+, SO42- Chemical properties of ions are different from the chemical properties of the atoms they come fromatom/ion are essentially the same, but the behaviors are quite different Predicting Ionic Charge Atoms will gain or lose electrons to have the same electronic structure as the noble gas closest to them- these electron arrangements are very stable and contribute to the nonreactivity of the noble gases Gaining/losing valence shell electrons to get to a full Periodic table is useful for keeping track of and predicting the charges of ions- atoms in groups have same number of valence electrons As a general trend, group 1A forms 1+ ions, 2A forms 2+ ions Group ions 6A forms 2- ions, and group 7A forms 1- Ionic compounds Ionic compound- compound that contains both positively and negatively charged ions Ex: NaCl Na reacts with Cl, Na will lose and electron, Cl will gain an electron Na+ will form, Cl- will form, and due to opposite charges, they attract one another (electrostatic) and form ionic bond Can predict whether a compound will be ionic (contain ions) or molecular (just molecules) Ionic compounds typically are formed between a metal and a nonmetal Molecular compounds- generally composed of nonmetals Ionic compounds arrange in a 3-D lattice structure- meaning they will alternate positive and negative charge (positive surrounded by negative, negative surrounded by positive) We can easily write the empirical formula of any ionic compound, as long as we know the charges of the ions Ions combine in ratios that allow the total positive charge to equal the total negative charge- ionic compounds combine to form a net zero charge Na+ will combine with Cl- to form NaCl (1and 1+ charge cancel out) If charges of ions are not equal, use subscripts to equal them out Ex: Ba2+ and Cl- will combine to form BaCl2 Sec. 2.8- Naming Inorganic Compounds Ionic Compounds For ionic Compounds 1. Cation Name (if multivalent, specify with roman numeral in parenthesis what charge. Ex: Copper (I) or Copper (II)) 2. Anion Name If not a polyatomic ion- ends in “-ide” If polyatomic, name of polyatomic (always ends in “-ite” or “-ate” No prefixes Acids For Acids: Anions end in “-ide”→ Add “Hydro-” prefix, change “-ide” to “-ic”, then add acid Anions end in “-ite” → Change “-ite” to “-ous”, then add acid Anions end in “-ate” → Change “-ate” to “-ic”, the add acid Binary Molecular Compounds Binary Molecular Compounds (two-element molecular compounds) 1. Name of element farthest to the left on periodic table written first (except compounds with Oxygen- oxygen not written first unless combined with F) 2. If both in same group- Higher atomic number named first 3. Name of second element ends in “-ide” 4. Greek prefixes- indicate number of atoms of each element. No “mono-” used for first element
