Water From Any Source Has Exactly the Same Hydrogen to Oxygen Ratio
Learn to verify the law of conservation of mass from reaction data, apply the law of constant proportions to find combining masses, and see which postulates of Dalton's theory explain each law.
Does water from a river and water from a laboratory have the same composition?
Exactly the same — and that is a law, not a coincidence.
Pure water from a river, from a well, from rainfall or made in a laboratory by burning hydrogen always contains hydrogen and oxygen in the **mass ratio **. So g of water contains g of hydrogen and g of oxygen, whatever its source.
That is the law of constant proportions, and it is one of two laws that govern how elements combine. The other is the law of conservation of mass, which says the total mass in a reaction never changes.
Both laws were established by careful weighing, and neither explains itself. What explains them is Dalton's atomic theory — and that relationship between a law and the theory behind it is the shape of this whole page. It covers the first part of the CBSE Class 9 Science chapter on the atomic foundations of matter.
Pure water from a river, from a well, from rainfall or made in a laboratory by burning hydrogen always contains hydrogen and oxygen in the **mass ratio **. So g of water contains g of hydrogen and g of oxygen, whatever its source.
That is the law of constant proportions, and it is one of two laws that govern how elements combine. The other is the law of conservation of mass, which says the total mass in a reaction never changes.
Both laws were established by careful weighing, and neither explains itself. What explains them is Dalton's atomic theory — and that relationship between a law and the theory behind it is the shape of this whole page. It covers the first part of the CBSE Class 9 Science chapter on the atomic foundations of matter.
How do you verify the law of conservation of mass from reaction data?
Add the masses of the reactants, add the masses of the products, and check that the two totals match.
The law states that mass is neither created nor destroyed in a chemical reaction — so the total mass of the reactants equals the total mass of the products.
Worked example 1 — a reaction with three products. g of sodium carbonate reacts with g of ethanoic acid, giving g of carbon dioxide, g of water and g of sodium ethanoate.
The two totals agree, so the law is verified for this reaction.
Worked example 2 — a decomposition. g of calcium carbonate is heated and gives g of calcium oxide and g of carbon dioxide.
Verified.
Worked example 3 — working out a missing mass. g of a metal combines with oxygen to form g of its oxide. How much oxygen reacted?
Worked example 4 — from atomic masses. Magnesium burns in oxygen to give magnesium oxide:
Taking Mg as u and O as u, , which is the mass of magnesium oxide. The equation balances by mass because the same atoms appear on both sides.
Why a burning candle seems to break the law. A candle plainly loses mass as it burns, and a rusting iron nail plainly gains some. Neither is a violation.
The candle's products — carbon dioxide and water vapour — escape into the air, so only part of the total is left on the plate. The nail takes oxygen from the air and adds it to itself. Weigh everything, including the gases, and the total is unchanged in both cases.
That is why the experiment must be done in a closed container. A sealed flask keeps every product inside, so the balance can weigh the whole system before and after. An open reaction cannot test the law, however carefully it is weighed — and recognising that is the point of the standard laboratory set-up.
The law states that mass is neither created nor destroyed in a chemical reaction — so the total mass of the reactants equals the total mass of the products.
Worked example 1 — a reaction with three products. g of sodium carbonate reacts with g of ethanoic acid, giving g of carbon dioxide, g of water and g of sodium ethanoate.
The two totals agree, so the law is verified for this reaction.
Worked example 2 — a decomposition. g of calcium carbonate is heated and gives g of calcium oxide and g of carbon dioxide.
Verified.
Worked example 3 — working out a missing mass. g of a metal combines with oxygen to form g of its oxide. How much oxygen reacted?
Worked example 4 — from atomic masses. Magnesium burns in oxygen to give magnesium oxide:
Taking Mg as u and O as u, , which is the mass of magnesium oxide. The equation balances by mass because the same atoms appear on both sides.
Why a burning candle seems to break the law. A candle plainly loses mass as it burns, and a rusting iron nail plainly gains some. Neither is a violation.
The candle's products — carbon dioxide and water vapour — escape into the air, so only part of the total is left on the plate. The nail takes oxygen from the air and adds it to itself. Weigh everything, including the gases, and the total is unchanged in both cases.
That is why the experiment must be done in a closed container. A sealed flask keeps every product inside, so the balance can weigh the whole system before and after. An open reaction cannot test the law, however carefully it is weighed — and recognising that is the point of the standard laboratory set-up.
How do you use the law of constant proportions to find combining masses?
Treat the fixed mass ratio as a proportion and scale it.
The law states that in a given compound, the elements are always present in a fixed ratio by mass, whatever the source of the compound or the method of preparing it.
Worked example 1 — water. Hydrogen and oxygen are in the ratio .
- g of water contains g of hydrogen and g of oxygen
- g contains g and g
- g contains g and g
Check against the formula. Water is , so one molecule holds u of hydrogen and u of oxygen:
The ratio from the formula matches the ratio from weighing, which is the whole reason formulas can be trusted.
Worked example 2 — carbon dioxide. Carbon and oxygen are in the ratio .
So g of carbon dioxide contains g of carbon and g of oxygen, and g contains g and g.
Worked example 3 — a limiting quantity. g of carbon is burnt in g of oxygen. What mass of carbon dioxide forms, and how much oxygen is left?
The ratio is fixed at , so g of carbon can combine with only g of oxygen:
The extra oxygen simply does not react. Supplying more of one element cannot change the ratio the compound forms in — that is exactly what the law forbids.
Worked example 4 — comparing two samples. Sample A of copper oxide contains g of copper and g of oxygen. Sample B contains g of copper and g of oxygen. Do they obey the law?
The same ratio, so both are the same compound and the law holds.
The law applies to a compound, not to a mixture. Air is a mixture, and its composition varies from place to place and with altitude. Water is a compound, and its composition never varies. So a question asking whether a substance obeys the law is often really asking whether it is a compound at all.
Two elements may form more than one compound. Carbon and oxygen form both carbon monoxide, with , and carbon dioxide, with . The law holds for each compound separately — it does not say a pair of elements can combine in only one ratio.
The law states that in a given compound, the elements are always present in a fixed ratio by mass, whatever the source of the compound or the method of preparing it.
Worked example 1 — water. Hydrogen and oxygen are in the ratio .
- g of water contains g of hydrogen and g of oxygen
- g contains g and g
- g contains g and g
Check against the formula. Water is , so one molecule holds u of hydrogen and u of oxygen:
The ratio from the formula matches the ratio from weighing, which is the whole reason formulas can be trusted.
Worked example 2 — carbon dioxide. Carbon and oxygen are in the ratio .
So g of carbon dioxide contains g of carbon and g of oxygen, and g contains g and g.
Worked example 3 — a limiting quantity. g of carbon is burnt in g of oxygen. What mass of carbon dioxide forms, and how much oxygen is left?
The ratio is fixed at , so g of carbon can combine with only g of oxygen:
The extra oxygen simply does not react. Supplying more of one element cannot change the ratio the compound forms in — that is exactly what the law forbids.
Worked example 4 — comparing two samples. Sample A of copper oxide contains g of copper and g of oxygen. Sample B contains g of copper and g of oxygen. Do they obey the law?
The same ratio, so both are the same compound and the law holds.
The law applies to a compound, not to a mixture. Air is a mixture, and its composition varies from place to place and with altitude. Water is a compound, and its composition never varies. So a question asking whether a substance obeys the law is often really asking whether it is a compound at all.
Two elements may form more than one compound. Carbon and oxygen form both carbon monoxide, with , and carbon dioxide, with . The law holds for each compound separately — it does not say a pair of elements can combine in only one ratio.
Which of Dalton's postulates explains each law?
Conservation of mass follows from atoms being indestructible; constant proportions follows from atoms being identical and combining in fixed numbers.
Dalton's postulates:
- All matter is made of very small particles called atoms
- Atoms are indivisible and can be neither created nor destroyed in a chemical reaction
- Atoms of a given element are identical in mass and chemical properties
- Atoms of different elements have different masses and different chemical properties
- Atoms combine in small whole-number ratios to form compounds
- The relative number and kinds of atoms are constant in a given compound
How conservation of mass is explained. A chemical reaction only rearranges atoms; none is created or destroyed. Since every atom keeps its own mass, the total mass of all the atoms cannot change. The sodium carbonate reaction in the earlier section gave g both before and after because the same atoms were present throughout, differently arranged.
How constant proportions is explained. If a compound always contains the same kinds of atoms in the same numbers, and atoms of an element always have the same mass, then the mass ratio has no freedom to vary. Water always has two hydrogen atoms to one oxygen atom, and hydrogen atoms always weigh u while oxygen atoms weigh u — so the ratio must be , every time.
Notice what the theory has done. The two laws were established by weighing, and they describe a pattern without saying why it holds. Dalton's postulates make both patterns follow from a picture of what matter is.
That is exactly the distinction drawn in the opening chapter of this course: a law describes, a theory explains. Here you can see one theory accounting for two separate laws, which is what gives a theory its reach.
The theory also makes a prediction. If atoms combine in small whole-number ratios, then the masses of an element combining with a fixed mass of another should be in simple whole-number ratios too. Carbon monoxide and carbon dioxide bear that out: for a fixed g of carbon, the oxygen masses are g and g, and is the simple ratio . A theory that predicts something it was not built to explain is a good theory — which is why Dalton's postulates were accepted so readily.
Dalton's postulates:
- All matter is made of very small particles called atoms
- Atoms are indivisible and can be neither created nor destroyed in a chemical reaction
- Atoms of a given element are identical in mass and chemical properties
- Atoms of different elements have different masses and different chemical properties
- Atoms combine in small whole-number ratios to form compounds
- The relative number and kinds of atoms are constant in a given compound
How conservation of mass is explained. A chemical reaction only rearranges atoms; none is created or destroyed. Since every atom keeps its own mass, the total mass of all the atoms cannot change. The sodium carbonate reaction in the earlier section gave g both before and after because the same atoms were present throughout, differently arranged.
How constant proportions is explained. If a compound always contains the same kinds of atoms in the same numbers, and atoms of an element always have the same mass, then the mass ratio has no freedom to vary. Water always has two hydrogen atoms to one oxygen atom, and hydrogen atoms always weigh u while oxygen atoms weigh u — so the ratio must be , every time.
Notice what the theory has done. The two laws were established by weighing, and they describe a pattern without saying why it holds. Dalton's postulates make both patterns follow from a picture of what matter is.
That is exactly the distinction drawn in the opening chapter of this course: a law describes, a theory explains. Here you can see one theory accounting for two separate laws, which is what gives a theory its reach.
The theory also makes a prediction. If atoms combine in small whole-number ratios, then the masses of an element combining with a fixed mass of another should be in simple whole-number ratios too. Carbon monoxide and carbon dioxide bear that out: for a fixed g of carbon, the oxygen masses are g and g, and is the simple ratio . A theory that predicts something it was not built to explain is a good theory — which is why Dalton's postulates were accepted so readily.
Where does Dalton's atomic theory break down?
On three of its postulates, and each failure points to a later part of the syllabus.
The atom is not indivisible. It contains electrons, protons and neutrons, as the previous chapter of this course established from the gold foil experiment. The word atom means uncuttable, and the name outlived the claim.
Atoms of the same element are not always identical in mass. Isotopes of one element have the same number of protons and different numbers of neutrons. Chlorine has two common forms, of mass numbers and — which is exactly why its average atomic mass is u and not a whole number.
Atoms of different elements do not always have different masses. Isobars are atoms of different elements with the same mass number. Calcium with and argon with both have a mass number of .
Three more limitations worth knowing:
- The theory does not explain why atoms combine at all. It states that they do, in fixed ratios, and offers no mechanism — which is precisely what chemical bonding supplies in the next part of this chapter.
- It does not explain allotropes. Diamond and graphite are both made of nothing but carbon atoms, and their properties are completely different. Identical atoms alone cannot account for that; the arrangement matters.
- Not every compound has a small whole-number ratio. Complex compounds such as sugar, , combine in ratios that are whole numbers but not small ones.
Being superseded is not the same as being wrong. Dalton's postulates still explain the two laws of chemical combination perfectly, and the arithmetic built on them — combining masses, formulas, molecular masses — is used unchanged today. What later work did was narrow the range in which the postulates hold, not discard them.
Each failure opens the next topic. The divisible atom leads to subatomic particles and electronic configuration. The unexplained combining leads to ionic and covalent bonding. The allotropes lead to structure. So the limitations of this theory are not a list to memorise for one question — they are the syllabus's route map for everything that follows.
The atom is not indivisible. It contains electrons, protons and neutrons, as the previous chapter of this course established from the gold foil experiment. The word atom means uncuttable, and the name outlived the claim.
Atoms of the same element are not always identical in mass. Isotopes of one element have the same number of protons and different numbers of neutrons. Chlorine has two common forms, of mass numbers and — which is exactly why its average atomic mass is u and not a whole number.
Atoms of different elements do not always have different masses. Isobars are atoms of different elements with the same mass number. Calcium with and argon with both have a mass number of .
Three more limitations worth knowing:
- The theory does not explain why atoms combine at all. It states that they do, in fixed ratios, and offers no mechanism — which is precisely what chemical bonding supplies in the next part of this chapter.
- It does not explain allotropes. Diamond and graphite are both made of nothing but carbon atoms, and their properties are completely different. Identical atoms alone cannot account for that; the arrangement matters.
- Not every compound has a small whole-number ratio. Complex compounds such as sugar, , combine in ratios that are whole numbers but not small ones.
Being superseded is not the same as being wrong. Dalton's postulates still explain the two laws of chemical combination perfectly, and the arithmetic built on them — combining masses, formulas, molecular masses — is used unchanged today. What later work did was narrow the range in which the postulates hold, not discard them.
Each failure opens the next topic. The divisible atom leads to subatomic particles and electronic configuration. The unexplained combining leads to ionic and covalent bonding. The allotropes lead to structure. So the limitations of this theory are not a list to memorise for one question — they are the syllabus's route map for everything that follows.
Exam tip
Exam tip: add both sides, and scale the ratio rather than guessing
For conservation of mass, total the reactants and total the products separately and state that they are equal. and .
Say why an open reaction seems to break the law — escaping gas for a candle, absorbed oxygen for a rusting nail — and that a closed container is needed to test it.
For constant proportions, write the fixed ratio first, then scale it. Water is , carbon dioxide , ammonia .
Check the ratio against the formula: gives .
In a limiting quantity question, only the ratio's worth of the excess reactant reacts: g of carbon uses g of oxygen and leaves the rest untouched.
To test whether two samples are the same compound, divide and compare: .
The law applies to a compound, not a mixture — air varies, water does not. And two elements may form several compounds, each with its own fixed ratio.
Name the postulate that explains each law: indestructible atoms for conservation of mass, identical atoms in fixed numbers for constant proportions.
List the limitations with their reasons: the atom is divisible, isotopes break identical mass, isobars break different mass, and allotropes and bonding go unexplained.
And give units on every mass — grams for laboratory quantities, u for atomic masses.
Say why an open reaction seems to break the law — escaping gas for a candle, absorbed oxygen for a rusting nail — and that a closed container is needed to test it.
For constant proportions, write the fixed ratio first, then scale it. Water is , carbon dioxide , ammonia .
Check the ratio against the formula: gives .
In a limiting quantity question, only the ratio's worth of the excess reactant reacts: g of carbon uses g of oxygen and leaves the rest untouched.
To test whether two samples are the same compound, divide and compare: .
The law applies to a compound, not a mixture — air varies, water does not. And two elements may form several compounds, each with its own fixed ratio.
Name the postulate that explains each law: indestructible atoms for conservation of mass, identical atoms in fixed numbers for constant proportions.
List the limitations with their reasons: the atom is divisible, isotopes break identical mass, isobars break different mass, and allotropes and bonding go unexplained.
And give units on every mass — grams for laboratory quantities, u for atomic masses.
Did you know
Why a rusting nail gets heavier and a candle gets lighter
Two everyday observations appear to contradict the law of conservation of mass in opposite directions, and both turn out to support it.
Leave an iron nail in damp air and it rusts. Weigh it before and after and the rusted nail is heavier. Nothing was added to it by hand, so where did the extra mass come from?
From the air. Rust is iron combined with oxygen and water, and both came from the atmosphere. Weigh the nail and the air around it as one closed system and the total is unchanged — the mass simply moved from the air into the nail.
Now burn a candle. It plainly gets lighter, and eventually almost nothing is left. Yet burning combines the wax with oxygen, so the products should weigh more than the wax did.
They do. The products are carbon dioxide and water vapour, and both are gases that drift away invisibly. Collect them and their combined mass exceeds the mass of wax that disappeared — by exactly the mass of oxygen consumed.
Both cases fool the eye for the same reason: gases are easy to overlook. A balance placed on a table weighs only what sits on the pan, and any reactant or product that arrives from the air or leaves into it is missed.
That is why the standard demonstration seals the reaction in a flask. Once nothing can enter or leave, the balance sees the whole system, and the reading before the reaction matches the reading after to the last decimal place.
So the law was not discovered by watching candles. It was discovered by weighing in closed vessels, and the two familiar observations above are examples of why the closed vessel matters rather than exceptions to the rule.
Leave an iron nail in damp air and it rusts. Weigh it before and after and the rusted nail is heavier. Nothing was added to it by hand, so where did the extra mass come from?
From the air. Rust is iron combined with oxygen and water, and both came from the atmosphere. Weigh the nail and the air around it as one closed system and the total is unchanged — the mass simply moved from the air into the nail.
Now burn a candle. It plainly gets lighter, and eventually almost nothing is left. Yet burning combines the wax with oxygen, so the products should weigh more than the wax did.
They do. The products are carbon dioxide and water vapour, and both are gases that drift away invisibly. Collect them and their combined mass exceeds the mass of wax that disappeared — by exactly the mass of oxygen consumed.
Both cases fool the eye for the same reason: gases are easy to overlook. A balance placed on a table weighs only what sits on the pan, and any reactant or product that arrives from the air or leaves into it is missed.
That is why the standard demonstration seals the reaction in a flask. Once nothing can enter or leave, the balance sees the whole system, and the reading before the reaction matches the reading after to the last decimal place.
So the law was not discovered by watching candles. It was discovered by weighing in closed vessels, and the two familiar observations above are examples of why the closed vessel matters rather than exceptions to the rule.
Exam relevance
How do the laws of chemical combination feed into JEE and NEET?
Because every quantitative calculation in chemistry begins from these two laws, and this page is where they are first used on numbers.
This is the foundation for the Class 11 Chemistry chapter Some Basic Concepts of Chemistry, examined in both JEE Main and NEET. That chapter opens with the same two laws, adds the law of multiple proportions and the law of reciprocal proportions, and then builds the mole concept on them. The prediction noted above — that oxygen masses combining with a fixed mass of carbon are in the ratio for the two oxides — is the law of multiple proportions stated in advance.
Where the arithmetic is reused. Stoichiometry is conservation of mass applied to a balanced equation, and every limiting-reagent problem is the * g of carbon in g of oxygen* calculation on this page with harder numbers. Percentage composition and empirical formula questions are the law of constant proportions run in reverse: given the masses, find the ratio of atoms.
Dalton's limitations connect forward to Class 11 Structure of Atom, where subatomic particles and isotopes are treated fully, and to Class 11 Chemical Bonding, which supplies the mechanism Dalton's theory lacked.
What the questions look like. Numericals dominate, and the standard shapes are verify the law from given masses, find the mass of one element in a stated mass of compound, and find how much of an excess reactant is left over. Assertion-reason items favour the statements that a rusting nail does not violate conservation of mass, and that the law of constant proportions applies to compounds and not mixtures. Match-the-column questions pair a law with the postulate that explains it.
How board and competitive emphasis differ. A board paper asks you to state both laws, list Dalton's postulates, and verify a law from one set of masses. A competitive paper sets a limiting reagent problem, or gives two compounds of the same elements and asks for the ratio of the combining masses — so the reasoning about which reactant runs out first matters more than the statements of the laws.
The single trap that costs the most marks. Assuming that adding more of one reactant produces more compound. It does not — the ratio is fixed, so the excess simply remains unreacted. In worked example 3 above, g of oxygen produced no more carbon dioxide than g would have. That idea becomes the limiting reagent in Class 11, and it is worth getting right here, where the numbers are simple.
A second trap worth naming. Attaching the wrong limitation to Dalton's theory. The atom being divisible and isotopes breaking the identical-mass postulate are two separate failures, and a question asking which postulate isotopes contradict wants the third one — atoms of a given element are identical in mass — and not the indivisibility postulate.
This is the foundation for the Class 11 Chemistry chapter Some Basic Concepts of Chemistry, examined in both JEE Main and NEET. That chapter opens with the same two laws, adds the law of multiple proportions and the law of reciprocal proportions, and then builds the mole concept on them. The prediction noted above — that oxygen masses combining with a fixed mass of carbon are in the ratio for the two oxides — is the law of multiple proportions stated in advance.
Where the arithmetic is reused. Stoichiometry is conservation of mass applied to a balanced equation, and every limiting-reagent problem is the * g of carbon in g of oxygen* calculation on this page with harder numbers. Percentage composition and empirical formula questions are the law of constant proportions run in reverse: given the masses, find the ratio of atoms.
Dalton's limitations connect forward to Class 11 Structure of Atom, where subatomic particles and isotopes are treated fully, and to Class 11 Chemical Bonding, which supplies the mechanism Dalton's theory lacked.
What the questions look like. Numericals dominate, and the standard shapes are verify the law from given masses, find the mass of one element in a stated mass of compound, and find how much of an excess reactant is left over. Assertion-reason items favour the statements that a rusting nail does not violate conservation of mass, and that the law of constant proportions applies to compounds and not mixtures. Match-the-column questions pair a law with the postulate that explains it.
How board and competitive emphasis differ. A board paper asks you to state both laws, list Dalton's postulates, and verify a law from one set of masses. A competitive paper sets a limiting reagent problem, or gives two compounds of the same elements and asks for the ratio of the combining masses — so the reasoning about which reactant runs out first matters more than the statements of the laws.
The single trap that costs the most marks. Assuming that adding more of one reactant produces more compound. It does not — the ratio is fixed, so the excess simply remains unreacted. In worked example 3 above, g of oxygen produced no more carbon dioxide than g would have. That idea becomes the limiting reagent in Class 11, and it is worth getting right here, where the numbers are simple.
A second trap worth naming. Attaching the wrong limitation to Dalton's theory. The atom being divisible and isotopes breaking the identical-mass postulate are two separate failures, and a question asking which postulate isotopes contradict wants the third one — atoms of a given element are identical in mass — and not the indivisibility postulate.
Key takeaways
The laws of chemical combination and Dalton's theory: quick revision
- Law of conservation of mass: mass is neither created nor destroyed in a chemical reaction, so total reactants total products.
- g sodium carbonate with g ethanoic acid gives g — matching g of reactants.
- g of calcium carbonate gives g. An g metal forming a g oxide used g of oxygen.
- Magnesium: u, which is the mass of magnesium oxide.
- A burning candle and a rusting nail do not break the law — gases escape in one case and are absorbed in the other. A closed container is needed to test it.
- Law of constant proportions: a given compound always has its elements in a fixed mass ratio, whatever its source.
- Water is , so g has g and g; g has g and g. The formula gives .
- Carbon dioxide is , so g has g of carbon and g of oxygen.
- Excess does not react: g of carbon in g of oxygen gives g of carbon dioxide and leaves g of oxygen.
- Two samples with and both give , so they are the same compound.
- The law applies to compounds, not mixtures — air varies, water does not. And two elements may form several compounds, each with its own ratio.
- Dalton's postulates: matter is made of atoms; atoms are indivisible and indestructible; atoms of one element are identical; atoms of different elements differ; they combine in small whole-number ratios; the number and kinds of atoms in a compound are constant.
- Conservation of mass is explained by atoms being merely rearranged; constant proportions by identical atoms combining in fixed numbers.
- The theory also predicts the oxygen ratio between carbon monoxide and carbon dioxide.
- Limitations: the atom is divisible; isotopes break the identical-mass postulate; isobars break the different-mass postulate; bonding and allotropes go unexplained; and sugar, , has a whole-number but not small ratio.
- A law describes and a theory explains — one theory here accounts for two laws.
Take any reaction whose masses you know and verify both laws from the same data — if the totals match and the ratio holds, you have used the laws rather than recited them.
- g sodium carbonate with g ethanoic acid gives g — matching g of reactants.
- g of calcium carbonate gives g. An g metal forming a g oxide used g of oxygen.
- Magnesium: u, which is the mass of magnesium oxide.
- A burning candle and a rusting nail do not break the law — gases escape in one case and are absorbed in the other. A closed container is needed to test it.
- Law of constant proportions: a given compound always has its elements in a fixed mass ratio, whatever its source.
- Water is , so g has g and g; g has g and g. The formula gives .
- Carbon dioxide is , so g has g of carbon and g of oxygen.
- Excess does not react: g of carbon in g of oxygen gives g of carbon dioxide and leaves g of oxygen.
- Two samples with and both give , so they are the same compound.
- The law applies to compounds, not mixtures — air varies, water does not. And two elements may form several compounds, each with its own ratio.
- Dalton's postulates: matter is made of atoms; atoms are indivisible and indestructible; atoms of one element are identical; atoms of different elements differ; they combine in small whole-number ratios; the number and kinds of atoms in a compound are constant.
- Conservation of mass is explained by atoms being merely rearranged; constant proportions by identical atoms combining in fixed numbers.
- The theory also predicts the oxygen ratio between carbon monoxide and carbon dioxide.
- Limitations: the atom is divisible; isotopes break the identical-mass postulate; isobars break the different-mass postulate; bonding and allotropes go unexplained; and sugar, , has a whole-number but not small ratio.
- A law describes and a theory explains — one theory here accounts for two laws.
Take any reaction whose masses you know and verify both laws from the same data — if the totals match and the ratio holds, you have used the laws rather than recited them.