The Building Blocks of Matter Chapter 3 - Atoms: The Building Blocks of Matter

From Philosophy to Science There were two schools of thought of the composition of the cosmos… is everything in the universe continuous and infinitely divisible Or, is there a limit to how small you can get? Particle theory was not the most popular early opinion, but was supported as early as Democritus in ancient Greece. From Philosophy to Science

From Philosophy to Science Democritus proposed that all the matter is composed of tiny particles called “Atomos” These “particles” were thought to be indivisible Aristotle did not accept Democritus’ atom, he was of the “matter is continuous” philosophy Because of Aristotle’s popularity his theory was adopted as the standard From Philosophy to Science

From Philosophy to Science By the 1700’s nearly all chemists had accepted the modern definition of an element as a particle that is indivisible It was also understood at that time that elements combine to form compounds that are different in their properties than the elements that composed them However, these understandings were based on observations not empirical evidence From Philosophy to Science

From Philosophy to Science There was controversy as to whether elements always combine in the same proportion when forming a particular compound. In the 1790’s, chemistry was revolutionized by a new emphasis on quantitative analysis because of new and improved balances This new technology led to the discovery of some new scientific understandings From Philosophy to Science

From Philosophy to Science The Law of Conservation of Mass: States that mass is neither created nor destroyed during ordinary chemical reactions or physical changes. Which means the total mass of the reactants must equal the total mass of the products. From Philosophy to Science

Law of Conservation of Mass +  Carbon, C Oxygen, O Carbon Monoxide, CO Mass x Mass y Mass x + Mass y Carbon Monoxide, CO Mass x + Mass y + Carbon, C Oxygen, O Mass x Mass y 

From Philosophy to Science The Law of Definite Proportions: The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or the source of the compound NaCl is NaCl no matter if it is table salt (small crystals) or rock salt (large crystals) From Philosophy to Science

From Philosophy to Science The Law of Multiple Proportions: If 2 or more different compounds are composed of the same 2 elements, then the ratio of the masses of the 2nd element combined with a certain mass of the 1st element is always a ratio of small whole numbers From Philosophy to Science

Law of Multiple Proportions + = Carbon Oxygen Carbon Monoxide, 1:1 1 1 + Carbon Oxygen = Carbon Dioxide, 1:2 1 2

In 1808, John Dalton proposed an explanation for each of the proposed laws He reasoned that elements were composed of atoms & that only whole #’s of atoms can combine to form compounds His ideas are now called the Atomic Theory of Matter Atomic Theory

All matter is composed of extremely small particles called atoms. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, & other properties. ELEMENT 1 ELEMENT 2 ELEMENT 3 ELEMENT 4 Atomic Theory

Atomic Theory Atoms cannot be subdivided, created, or destroyed atoms of different elements combine in simple whole # ratios to form chem compds in chemical rxns, atoms are combined, separated, or rearranged + + Atomic Theory

Through these statements, evidence could be gathered to confirm or discount its claims Not all of Dalton’s claims held up to the scrutiny of experimentation Atoms CAN be divided into even smaller particles Not every atom of an element has an identical mass Atomic Theory

Atomic Theory Dalton’s Atomic Theory of Matter has been modified. What remains unchanged is… All matter is composed of atoms Atoms of any one element differ in properties from atoms of another element Atomic Theory

One of the disputed statements of Dalton was that atoms are indivisible. In the 1800’s it was determined that atoms are actually composed of several basic types of smaller particles it’s the number and arrangement of these particles that determine the atom’s chemical properties. Atomic Theory

The definition of an atom that emerged was: the smallest particle of an element that retains the chemical properties of that original element. All atoms consist of 2 regions that contain the subatomic particles The nucleus The electron cloud around the nucleus Atomic Theory

The nucleus is a very small region located near the center of the atom In every atom the nucleus contains at least 1 proton, which is positively charged particle and usually contains 1 or more neutral particles called neutrons (Hydrogen does not contain a neutron.) Atomic Structure

The electron cloud is the region that surrounds the nucleus This region contains 1 or more electrons, which are negatively charged subatomic particles The volume of the electron cloud is much larger than the nucleus Atomic Structure

With the exception of Hydrogen, every nucleus contains 2 kinds of particles protons and neutrons they make up the mass of the atom (Mass Number = Protons + Neutrons) Proton has a charge equal to but opposite of the charge of an electron. Atoms are neutral because they contain equal #’s of protons & electrons Atomic Structure

The atoms of different elements differ in the # of protons in their nuclei and therefore in their positive charge The # of protons the atom contains determines the atom’s identity, also known as atomic number. Only Oxygen contains 8 protons Only Fluorine contains 9 protons Only Neon contains 10 protons Structure of the Atom

Particle Symbol Charge Mass Number Relative Mass (amu) Actual Mass (kg) Electron e- -1 0.0005486 9.109 x 10-31 Proton p+ +1 1 1.007276 1.673 x 10-27 Neutron n0 1.008665 1.675 x 10-27

Ch 3.3: Atomic Number

# protons in an atom = # electrons Elements are identified by the number of PROTONS they contain. The “atomic number” of an element is the number of protons in the nucleus PROTONS IDENTIFIES AN ELEMENT!!! # protons in an atom = # electrons Why? Because atoms are neutral!

Complete Symbol X Mass number Atomic Subscript → Superscript →

Cl 35 17 # OF PROTONS + # OF NEUTRONS MASS NUMBER ATOMIC NUMBER NUMBER OF PROTONS # OF PROTONS + # OF NEUTRONS

Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope: Mass # = p+ + n0 Element p+ n0 e- Mass # Oxygen - 10 - 33 42 - 31 15 Mass number 18 8 8 18 Arsenic 75 33 75 Phosphorus 16 15 31

Find the # of e-, p+ and n0 for sodium. (mass # = 23) Practice Problems Find the # of e-, p+ and n0 for sodium. (mass # = 23) Find the # of e-, p+ and n0 for uranium. (mass # = 238) Atomic # = 11 = # e- = # p+ # neutrons = 23-11 = 12 Atomic # = 92 = # e- = # p+ # neutrons = 238-92 = 146

Check for understanding: If an element has 91 protons and 140 neutrons find the: Atomic number Mass number number of electrons element name 91 231 91 protactinium

(Isotopes: The # after the name is the mass #.) An isotope refers to atoms that have the same # of protons, but a different number of neutrons. Because of this, they have different mass #’s. Ex---> (1) Carbon-12 & Carbon-13 (2) Chlorine-35 & Chlorine-37 (Isotopes: The # after the name is the mass #.)

Cl Cl 18 20 35 37 17 17 EXAMPLE OF AN ISOTOPE ATOMIC MASS ATOMIC NUMBER 18 NEUTRONS 20 NEUTRONS

Br Question #1 80 35 Find each of these: Atomic number Mass Number number of protons number of neutrons number of electrons Br 35

Question #2 If an element has an atomic number of 34 and a mass number of 78, what is the: number of protons number of neutrons number of electrons complete symbol

Units = atomic mass unit (amu) 12 Units = atomic mass unit (amu) The atomic masses listed in the Periodic Table are a “weighted average” of all the isotopes of the element.

In chemistry, chlorine has 2 isotopes: Weighted Average Practice Problems: In chemistry, chlorine has 2 isotopes: Cl-35 (75.8% abundance) Cl-37 (24.23 % abundance) What is the weighted average atomic mass of chlorine? 35 x 0.758 = 26.53 amu 37 x 0.2423 = 8.965 amu Add them up!!! This rounds to 35.5 amu + 35.495 amu

Relating Mass Numbers to Atoms The Mole: the amount of a substance that contains as many particles as there are atoms in exactly 12 grams of carbon-12. Avogadro’s Number: the number of particles in exactly one mole of a pure substance = 6.022 x 1023. Molar Mass: the mass of one mole of a pure substance. Units = g/mol

This is when we get to use dimensional analysis! The conversion factors we need are: and of course…molar mass

Gram to Mole Conversions Mass of Element in Grams Number of Moles of Element Number of Atoms of Element

Practice Problem ALWAYS USE PARANTHESES AROUND YOUR CONVERSION FACTORS!! You have 3.50 mol of Copper. What is it mass in grams? You have 28.55 grams of Carbon. How many atoms is this? You have 4.85 x 1030 atoms of Magnesium. How many moles is this?