Chemistry Counting Particles too small to see. John Dalton (1776-1844) Proposed Atomic Theory Each element is made up of tiny, indivisible atoms. Different.

Slides:



Advertisements
Similar presentations
Stoichiometry! The math of chemistry .
Advertisements

Atomic Theory and Chemical Reactions SEPUP Science in Global Issues Chemistry: Fueling the World Activity 12 Field Test (Spring, 2008) © 2008 Regents of.
Chapter 10: Chemical Quantities
The Mole – A measurement of matter
Amedeo Avogadro In order to state that equal volumes of all gases at the same temperature and pressure contain the same number of molecules, Amedeo Avogadro.
Section 10.1 Measuring Matter
Lecture No. 1 Laws of Chemical Combinations Chemistry.
Molecular Composition of Gases Volume-Mass Relationships of Gases.
Chapter 7 Chemical Quantities or How do you measure how much? You can measure mass, volume, or you can count pieces of a substance. We measure mass in.
A NOTHER W AY TO M EASURE G ASES - V OLUME. The volume of a gas can be influenced by both temperature and pressure. Because of this behavior, it is necessary.
Chemical Quantities Math in Chemistry. Measuring Matter measure the amount of something by one of three different methods— by count, by mass, and by volume.
Chemistry Chemical Reactions – Rearranging Atoms.
Chapter 10 & 11 Chemical quantities and Chemical Reactions.
Chapter 11: Molecular Composition of Gases
Empirical and Molecular Formulas
Ch. 11 Molecular Composition of Gases
Preview Lesson Starter Objectives Measuring and Comparing the Volumes of Reacting GasesMeasuring and Comparing the Volumes of Reacting Gases Avogadro’s.
Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Main AR Standards.
Unit 10 – The Mole Essential Questions:
1 Chapter 10 “Chemical Quantities” Chemistry Pioneer High School Mr. David Norton.
The Mole Concept. Relative Mass The relative mass of an object is the mass of that object as a multiple of some other object’s mass. In the example, the.
The Mole 1 dozen = 12 1 gross = ream = mole = 6.02 x 1023.
Stoichiometry By Ellis Benjamin. Definitions I Compounds - is a pure substance that is composed of two or more elements Molecules – is a combination of.
1 Chemistry 11 Chapter 4 - The MOLE. 2 Relative Atomic Mass Dalton, concerned with how much one element could combine with a given amount of element,
Unit 12 Mass and Moles.
Counting Atoms Chapter 9. MOLE?? Moles of Particles In one mole of a substance, there are 6 x particles.
Chapter 7 Chemical Quantities Spring The Mole: A Measurement of Matter- What Is a Mole?  We use problem solving steps to figure out the amount.
CHAPTER 10 THE MOLE. The mole is a number (6.02 x ) It is a term like the term “dozen” It was chosen by chemists to make working with atomic weights.
The Mole Chapter 11 – Chemistry L1 LSM High School Section 11.1: Measuring Matter Objectives: Describe how a mole is used in chemistry Relate a mole to.
The Mole Chemistry 6.0 The Mole I. Formulas & Chemical Measurements A. Atomic Mass 1. Definition: the mass of an atom, based on a C-12 atom, in atomic.
Chapter 10 The Mole. Chemical Measurements Atomic Mass Units (amu) – The mass of 1 atom – 1 oxygen atom has a mass of 16 amu Formula Mass (amu or fu)
General Chemistry Chapter 2 Definitions Left click your mouse to continue.
1 Chapter 10 “Chemical Quantities” Yes, you will need a calculator for this chapter!
THE MOLE. Atomic and molecular mass Masses of atoms, molecules, and formula units are given in amu (atomic mass units). Example: Sodium chloride: (22.99.
Chapter 7 Chemical Quantities Fall The Mole: A Measurement of Matter- What Is a Mole?  How do you measure matter?  You count things  You weigh.
Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Main AR Standards.
MOLES!! Chemical Quantities. Counting by weighing  When things are too small to count out individually, we can “count” them by weighing them. Must know.
Quantities in Chemistry The Mole and Molar Mass. Mole Review A Mole is a unit of measurement in chemistry. It represents 6.02 x of an entity. One.
Atomic Theory The Mole. Things that you already know Science measures a lot of things. Examples of what we measure include time, mass, temperature, volume,
Starter S-89 1.List the elements and how many there are of each in Iron (III) oxide. 2.Multiply the number of each element by the average atomic mass of.
Gas Laws Joseph Louis Gay-LussacAmadeo Avogadro Robert BoyleJacques Charles.
1. 2 Chemical Quantities or 3 How you measure how much? How you measure how much? n You can measure mass, n or volume, n or you can count pieces. n We.
Law of conservation of mass states that mass can neither be created nor destroyed in a chemical reaction.
Preview Lesson Starter Objectives Measuring and Comparing the Volumes of Reacting GasesMeasuring and Comparing the Volumes of Reacting Gases Avogadro’s.
Stoichiometry and the mole Chapter 8 What is stoichiometry?  Quantitative aspects of chemistry  Stoicheon Greek root (element)  Metron Greek root(
Unit 5 Review The Mole. 1. What is the mass of 1 mole of iron atoms? A amu B L C x D g.
Chemical Composition … Moles Chemistry Mr. Lentz.
Unit 6 Review The Mole.
Chapter 7 Chemical Quantities or How you measure how much? You can measure mass, volume, or you can count pieces of a substance. We measure mass in grams.
The Mole  Just to clear up any misconceptions when we use the term “mole” we are not referring to this small blind fellow.
The Mole Intro to Stoichiometry. Measurements in Chemistry Atomic Mass: the mass of an atom of a certain element in atomic mass units (amu). 1 amu = 1.66.
Miss Fogg Spring 2016  A particle can refer to an individual atom OR a type of molecule ◦ Jellybean ◦ Baseball ◦ Carbon atoms ◦ Hydrogen atoms ◦ Water.
Gas volumes and moles PAGE 87 OF INB. Essential Question:  How can 2 liters of Hydrogen react with 1 liter of Oxygen and only produce 2 liters of gas?
The Mole. REVIEW OF TERMINOLOGY Atomic mass mass of 1 atom in AMUs read from P.T. Atomic Mass C = AMU H = AMU.
Unit V: The Mole Concept
Chemistry10.1.
Mass Relationships and Avogadro’s Number
Chapter 10.1 The Mole: A Measurement of Matter
Chemical Quantities.
Quantitative chemistry
Reactions of Gases Lesson 8.
Chapter 10 – Chemical Quantities
Ch. 11: Molecular Composition of Gases
Unit 2: Chemistry Lesson 2: Classifying Matter Essential Questions: 1
The Mole.
Unit 2: Chemistry Lesson 2: Classifying Matter Essential Questions: 1
Created by C. Ippolito June 2007
Chapter 11 Preview Lesson Starter Objectives
Chapter 11 Gas Volumes and the Ideal Gas Law Section 3.
Starter S-93 What is the molar mass of H2S?
Presentation transcript:

Chemistry Counting Particles too small to see

John Dalton ( ) Proposed Atomic Theory Each element is made up of tiny, indivisible atoms. Different elements have different atoms. Compounded substances (compounds) are composed of fixed groupings of atoms called molecules (“little lumps of matter”)

Joseph Louis Gay-Lussac ( ) Discovered that gases at constant pressure and temperature always combine in definite volume ratios. 1 Liter of hydrogen gas + 1 Liter of chlorine gas always produced 2 Liters of hydrogen (mono)chloride. 2 Liters of hydrogen gas + 1 Liter of oxygen gas always produced 2 Liters of water vapor

Joseph Louis Gay-Lussac ( ) Concluded equal volumes of gas must have the same number of particles. 1 Liter of hydrogen gas has the same number of hydrogen particles as 1 Liter of oxygen gas. 2 Liters of water vapor has 2X the number of particles as 1 Liter of oxygen gas.

Consequences of Gay- Lussac’s Findings How is it possible to get two volumes (liters) of product when you only started with one volume (liter) of reactants? But the results continued to yield two volumes of gas C Instead of one Should Equal 5 atoms of gas A 5 atoms of gas B 5 molecules of gas C

Amedeo Avogadro ( ) Explained Lussac’s findings by: Realizing that some gas particles are actually diatomic molecules instead of atoms. Diatomic means they are made up of two of the same atoms bonded together. Equal 5 molecules of gas A 5 molecules of gas B 10 molecules of gas C

Empirical Formula Based on the ratio of volumes of reactants and products, a chemical formula for the product can be deduced 2 volumes of hydrogen combine with 1 volume of oxygen to produce 2 volumes of water Thus, the formula for water must be H 2 O.

Relative Mass of Atoms Scientists also noticed that gases combined in specific mass proportions. Two volumes of hydrogen gas (each with a mass of 0.5g) always combined with one volume of oxygen gas (with a mass of 8g) Thus, oxygen’s mass relative to hydrogen is

This means: 1 volume of oxygen has the same number of particles as 1 volume of hydrogen, but the mass of the volume oxygen is 16 times greater than the mass of the volume of hydrogen. If the relative mass of hydrogen is 1, then the relative mass of oxygen would be 16. Relative Mass of Atoms

The Mole Weighable amount of an element or compound. If the relative mass of hydrogen is 1 and the relative mass of oxygen is 16, then H 2 O molecule would have a relative mass of 18. Relative mass of 1 Relative mass of 1 Relative mass of 16

Molecular Mass The mass of the smallest unit of an element or compound. Determined by the mass of the elements in the formula and the number of atoms. Measured in Atomic Mass Units (amu) because it is more convenient than grams The molecular mass of H 2 O is 18 amu

Molar Mass The number of grams equivalent to the molecular mass of an element or compound. 1amu is equivalent to 1gram If the molecular mass of H 2 O is 18amu, then the molar mass of H 2 O is 18 grams. 1 Molar mass of a substance = 1 mole of substance.

Number of Particles in a Mole Thus, one molar mass (mole) of any substance (element or compound) has the same number of particles. 1 gram of substance = x amu. This means there are x particles of a substance in one mole of that substance.

1 mol of a substance = 1 molar mass 1molar mass = x particles The molar mass of the substance is equal to the molecular mass of the substance determined by the chemical formula for the substance. How Much?