Atomic Structure Chapter 4

The Early History of Chemistry The Greeks were the first to try to explain why chemical changes occur Aristotle (384-322 B.C.) proposed that matter was composed of four fundamental substances: fire, earth, water, and air Democritus (460-370 B.C.) proposed that matter was composed of small, indivisible particles Atomos, Greek for “uncuttable” Lacked experimental support Was not based on the scientific method No evidence to support either idea

Dalton’s Atomic Theory John Dalton (1766-1844), an English chemist and schoolteacher Wanted to find explanations for these three laws: Law of Conservation of Mass Law of Multiple Proportions Law of Definite Proportions

Three Laws 1.) Law of Conversation of Mass: During a chemical reaction, mass is neither created nor destroyed (mass of reactants equals the mass of the products). 2.) Law of Definite Proportions: A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. Example: Water (H2O) always consists of 89% (by mass) oxygen and 11% (by mass) hydrogen

Three Laws (cont.) Ratio of oxygen in NO2 to oxygen in NO is 2:1 3.) Law of Multiple Proportions: When two elements can combine to form more than one compound (i.e., CO and CO2), then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers Example: Ratio of oxygen in NO2 to oxygen in NO is 2:1

Postulates of Darwin’s Atomic Theory All matter is composed of extremely small, indivisible particles called atoms All atoms of a given element are identical (same properties); the atoms of different elements are different Atoms are neither created nor destroyed in chemical reactions, only rearranged 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds

Now We Know…Atoms are Divisible! Dalton’s atomic theory is accepted still today but with two important changes By the 1850s, scientists began to realize that the atom was made up of subatomic particles These subatomic particles are protons, neutrons, and electrons Also, atoms of the same element can be different (isotopes, which we will discuss later)

Modern Atomic Theory 1. All matter is composed of atoms 2. Atoms are the smallest unit of matter that retains its identity 3. Atoms come in over 115 basic types called elements 4. Atoms are very small (2.4 x 1022 atoms in a penny) 5. The mass of an atom is concentrated in its positively-charged nucleus

Early Models of Atomic Structure

J. J. Thomson (1897) Discovered the electron (negatively-charged particles in the atom) Devised the “Plum Pudding Model” of atomic structure

Plum Pudding Model Thompson understood that atoms are neutral, so there must be a balance to the electron’s negative charge Proposed that positively- charged particles and the negatively-charged electrons were spread evenly throughout the atom Experiments by other scientists proved this model to be incorrect

Ernest Rutherford (1911) Determined the placement of the positively-charged particles (protons) by his Gold Foil Experiments Proposed the existence of a positively-charged nucleus

Gold Foil Experiments Rutherford studied angles at which alpha particles (which are positively-charged) were scattered as they passed through a thin gold foil If the “Plum Pudding Model” was correct, the alpha particles should pass directly through the foil, with perhaps only a slight deflection For the majority of the particles, this is indeed what happened For other particles, they were deflected back towards the original source Rutherford concluded two things: Within the atom exists particles (now we know are positively- charged protons) with a force strong enough to repel the alpha particles The part of the atom where these particles are contained (now we know is the nucleus) must be very small, since very few of the alpha particles were deflected

Gold Foil Experiments “Plum Pudding Model”…what should have happened What actually happened…

Niels Bohr (1913) Proposed the “Planetary Model” to describe the placement of electrons The “Planetary Model” stated that the electrons orbit around the nucleus, much like how planets orbit around the sun

Planetary Model

Modern Model of Atomic Structure

Modern Model of the Atom Nucleus: Contains protons and neutrons, so it contains the majority of an atom’s mass; positively-charged due to the presence of the protons Electron cloud: Mostly empty space that is responsible for the majority of the volume of the atom; where the negatively-charged electrons can be found orbiting the nucleus

Atomic Mass Unit (amu) Example: Relative mass of atom, using Carbon-12 as the standard Proton and neutron masses are both close to 1 amu 1 amu = 1/12 of the mass of a Carbon-12 atom 6 protons (~6 amu) + 6 neutrons (~6 amu) = 12 amu Example: Hydrogen contains 1 proton (~1 amu) and 0 neutrons. 1 proton (~1 amu) + 0 neutrons (0 amu) = 1 amu

Major Subatomic Particles Name Symbol Charge Relative Mass (amu) Actual Mass (g) Electron e- -1 1/1840 9.11x10-28 Proton p+ +1 1 1.67x10-24 Neutron no Atoms are measured in picometers (which is 10-12 meters) Nucleus is even smaller Radius of the nucleus is approximately 10-15 m

Distinguishing Atoms of Different Elements

Describing Atoms of Different Elements Elements are different because they contain different numbers of protons Atomic Number (Z) = number of protons (p+) in the nucleus Determines the type of atom (i.e., Li atoms always have 3 protons in the nucleus, Hg always 80) Mass Number (A) = number of protons + number of neutrons Electrons have a negligible contribution to overall mass In a neutral atom, there is the same number of electrons (e-) and protons (atomic number)

E Z Nuclear Symbols A Mass Number Elemental Symbol Atomic Number Every element is given a corresponding nuclear or chemical symbol which is composed of 1 or 2 letters (first letter upper case, second lower), as well as the mass number (A; # of protons + # of neutrons) and atomic number (Z; # of protons) E A Z Elemental Symbol Mass Number Atomic Number

Isotopes Atoms of the same element can have different numbers of neutrons, and therefore have different mass numbers The atoms of the same element that differ in the number of neutrons are called isotopes of that element When naming, write the mass number after the name of the element

Isotopes of Hydrogen H 1 Hydrogen-1 2 Hydrogen-2 3 Hydrogen-3

Properties of Isotopes Chemical properties are primarily determined by the number of electrons All isotopes have the same number of electrons, so they have nearly identical chemical properties even though they have different masses HOWEVER, physical properties often depend on the mass of the particle, so isotopes will have slightly different physical properties such as density, boiling point, melting point, etc.

Practice Give the number of protons, neutrons, and electrons for each of the following atoms, along with each one’s nuclear symbol Carbon-12 Carbon-13 Carbon-14 Boron-10 Boron-11 Chlorine-35 Chlorine-37

How heavy is an atom of oxygen? There are different kinds of oxygen isotopes: 16O, 17O, 18O We are usually more concerned with an average atomic mass (the decimal number you see on the Periodic Table for each element) Atomic mass is based on relative abundance of each isotope found in nature 15.999 amu is the atomic mass of oxygen It is the weighted average of all the isotopes Measured in atomic mass units (amu)

Atomic Mass The atomic mass on the Periodic Table is not a whole number, because it is a weighted average mass value of all the isotopes of the element To calculate the atomic mass of an element, multiply the mass (in amu) of each isotope by its natural abundance (often expressed as a percentage; convert it to a decimal) and then add all the products

Calculating Averages Example: Average = (% as decimal) x (mass1) + (% as decimal) x (mass2) + (% as decimal) x (mass3) + … Example: Silver has two naturally occurring isotopes: 107Ag with a mass of 106.90509 amu and abundance of 51.84 % 109Ag with a mass of 108.90476 amu and abundance of 48.16% What is the atomic mass of silver?

Practice Calculate the atomic mass of copper if copper has two isotopes 69.1% has a mass of 62.93 amu The rest (30.9%) has a mass of 64.93 amu Magnesium has three isotopes 78.99% magnesium 24 with a mass of 23.9850 amu 10.00% magnesium 25 with a mass of 24.9858 amu The rest magnesium 26 with a mass of 25.9826 amu What is the atomic mass of magnesium?

Practice Calculate the atomic mass for Chlorine-35 and Chlorine-37 Calculate the atomic mass for Nitrogen-14 and Nitrogen-15 Calculate the atomic mass for Bromine-79 and Bromine-81 Cl-35 =75.78 Cl-37 = 24.22% N-14 = 99.632% N-15 = 0.368% Br-79 = 50.69% Br-81 = 49.31%